S34* 


IRATORY  EXPERIMENTS 


GENERAL  CHEMISTRY 


HMAN         SCHLUNDT 


UNIVERSITY  OF  CALIFORNIA 
AT   LOS  ANGELES 


GIFT  OF 

LECK 


LABORATORY  EXPERIMENTS 

IN 

GENERAL  CHEMISTRY 


By 
HERMAN  SCHLUNDT 

Professor  of  Physical  Chemistry 
University  of  Missouri 


SECOND  EDITION— REVISED 


COLUMBIA,  MISSOURI 

PRESS  OF  E.  W.  STEPHENS  PUBLISHING  COMPAHT 
1912 


Copyright,  1912,  by 
HERMAN  SCHLUNDT 


-;  CONTENTS 

7  CHAPTERS  TITLES                                                               PAGES 

General  Instructions 7 

I.     Apparatus 9 

'Y)         II.     Hydrogen 18 

*T)       III.     Oxygen 23 

^        IV.     Water 28 

V.     Equivalent  Weights,  Formulas,  Equations 31 

VI.     Review  Exercises 38 

VII.     The  Halogens 40 

^  VIII.     Acids,  Bases,  Salts.     Chemical  Equilibrium 47 

}      IX.     Ammonia  and  Nitric  Acid 52 

I        X.  Hydrogen  Sulphide,  Sulphur  Dioxide,  Carbon  Dioxide  55 

j      XI.  The  Atmosphere,  Flame.     Oxidation  and   Reduction . .  63 

v\    XII.     lonization 70 

rv^XIII.     Elective  Qualitative  and  Quantitative  Experiments 72 

XIV.     Some  Inorganic  and  Organic  Preparations 76 

';sL               Appendix 84 


I 


232445 


PREFACE 

The  experiments  outlined  in  this  manual  are  designed  primarily 
for  college  students  who  have  not  had  a  course  in  chemistry  in  a 
preparatory  school.  The  exercises  represent  the  laboratory  work 
of  a  comparatively  brief  introductory  course  in  General  Chemistry. 
The  experiments  are  to  be  conducted  under  the  guidance  of  an  in- 
structor, and  are  to  be  supplemented  by  class-room  demonstrations 
in  connection  with  recitations  from  a  text-book  in  General  Chemistry 
for  college  students,  or  by  illustrated  lectures  and  text-book  assign- 
ments. 

Frequently  the  experiments  do  not  furnish  sufficient  information 
to  enable  the  student  to  answer  some  of  the  questions  and  make  the 
explanations  that  are  to  appear  in  his  notebook.  The  necessary 
information  can  generally  be  obtained  from  the  text-book,  and  it  is 
my  plan  to  have  the  student  use  the  text-book  and  laboratory  outline 
as  companion  volumes  in  the  laboratory.  To  facilitate  the  stu- 
dent's progress  in  this  connection  page  references  to  two  widely  used 
texts  have  been  inserted.  My  experience  goes  to  show  that  the 
student  will  thus  make  very  efficient  use  of  his  time,  that  he  will  give 
care  and  thought  to  his  work,  and  that  the  laboratory  work  can  be 
successfully  made  the  central  feature  of  instruction  in  the  course. 

In  preparing  the  experiments,  the  substances  chosen  for  study 
have  purposely  been  limited  to  avoid  scattering  the  student's  efforts. 
Intensive,  rather  than  extensive,  study  has  been  the  underlying 
idea  in  sleeting  the  exercises.  Extended  experiments  on  the  metal- 
lic elements  have  not  been  included,  as  I  feel  that  this  work  should 
be  undertaken  in  Analytical  Chemistry,  and  be  allotted  some  of  the 
time  so  largely  used  for  laboratory  practice  in  following  a  scheme  of 
separations.  I  fully  realize  the  value  of  practice  in  the  identification 
of  "unknown"  substances,  and  this  feature  of  laboratory  work  has 
been  duly  emphasized,  and,  it  is  hoped,  in  a  manner  which  preserves 
its  educational  value.  Emphasis  has  also  been  placed  upon  the  gen- 
eral reactions  of  acids,  bases,  and  salts,  the  processes  of  oxidation 
and  reduction,  and  chemical  changes  prominent  in  everyday  life.  My 
experience  has  been  that  the  work  outlined  is  adequate  in  scope  as 

(4) 


Preface  5 

a  preparation  for  more  advanced  courses  in  chemistry,  and  that  it 
offers  the  cultural  benefits  of  laboratory  work  to  the  student  who 
wishes  to  take  only  elementary  chemistry  as  a  part  of  his  college 
course.  The  benefits  of  laboratory  work,  of  course,  lie  more  in  the 
hands  of  the  instructor  than  in  the  pages  of  the  book,  and  fully  as 
much  in  the  enthusiasm  and  spirit  of  inquiry  of  the  student  as  in 
text-book  matter  and  qualities  of  the  teacher.  In  the  pages  of  this 
little  book  I  have  sought  throughout  to  enforce  the  scientific  method 
of  work.  The  educational  value  of  the  experiments  will  be  realized 
by  the  student  in  so  far  as  he  succeeds  in  making  the  scientific  method 
a  life  habit. 

The  majority  of  the  experiments  have  been  collected  from  various 
sources  and  modified  in  some  respects  by  substitutions  and  additions 
of  original  material.  Helps  and  suggestions  to  the  student  have 
been  distributed  in  the  form  of  notes  with  a  view  of  encouraging 
him  in  his  efforts,  and  gradually  building  up  higher  standards  and 
ideals.  The  optional  experiments  are  designed  for  the  faster  workers 
in  the  class,  and  will  serve  a  very  useful  purpose  in  keeping  the  am- 
bitious student  from  getting  too  far  ahead  of  the  recitation  work  of 
the  class.  The  interest  of  the  bright,  quick  student  often  flags,  and 
the  quality  of  his  work  deteriorates  because  he  reads  ahead  of  the 
class,  and  then  fails  to  prepare  the  regular  assignment.  The  chapter 
of  elective  experiments  will  be  found  of  excellent  service  in  this  con- 
nection. 

I  am  indebted  to  the  Laboratory  Outline  of  Alexander  Smith  for 
many  ideas  in  laboratory  instruction.  I  also  desire  to  express  my 
thanks  to  Mr.  E.  E.  Morlan  of  this  University  for  helpful  suggestions 
in  the  selection  of  material  for  the  experiments.  Corrections  and 
suggestions  will  be  appreciated. 

HERMAN  SCHLUNDT. 
University  of  Missouri,  August,  1910. 


PREFACE  TO  THE  SECOND  EDITION 

The  content  and  order  of  experiments  in  the  first  edition  have  in 
the  main  been  retained:  Some  of  the  quantitative  experiments  have 
been  replaced  by  others  and  several  additional  quantitative  experi- 
ments have  been  introduced.  Minor  changes  have  been  made  in 
a  number  of  experiments,  and  more  of  the  experiments  have  been 
designated  for  two  students  working  together.  Experience  has 
shown  thafe  the  plan  of  having  students  work  in  pairs  at  times  en- 
courages discussion  and  materially  increases  interest  and  progress. 
I  am  indebted  to  our  assistants  and  instructors  for  many  of  the  changes 
introduced.  Suggestions  will  be  welcomed  from  those  that  may  use 
this  manual.  H.  S. 

University  of  Missouri,  December,  1911. 


(6) 


GENERAL  INSTRUCTIONS 

1.  Check  the  outfit  of  apparatus  listed  on  the  printed  card  in 
your  locker.     To  do  this,  put  all  the  articles  on  the  top  of  the  desk, 
then  check  and  return  such  pieces  as  you  can  identify.     Since  articles 
missing,  broken,  or  imperfect  will  be  charged  for  when  the  course  is 
completed,  each  piece  as  it  is  returned  to  the  locker  should  be  care- 
fully examined,  and  any  shortages  or  defective  apparatus  should  be 
reported  to  an  instructor.     Place  the  test  tubes  in  the  rack  provided 
for  them,  and  arrange  the  glassware  so  that  opening  and  closing  the 
drawers  will  not  break  or  damage  the  apparatus.     Any  unfamiliar 
articles  may  be  checked  with  the  aid  of  an  instructor.     Finally  sign 
the  card  and  hand  it  to  the  instructor. 

2.  Provide   yourself   with   a   notebook  and    make   a   careful   per- 
manent record  of  each  experiment  unless  the  directions  state  other- 
wise.    Enter  the  numbers  and  titles  of  the  experiments  as  they  ap- 
pear in  this  manual.     Record  what  you  observe  in  short,  clear  state- 
ments and  at  the  time  the  observations  are  made.      Having  made  a 
record    of   the   observed   facts,    enter  the  conclusions  you  draw.     Do 
not  copy  the  directions,  but  direct  your  efforts  to  an  intelligent  inter- 
pretation of  the  data  presented  by  the  experiment,   and  then  express 
your  ideas  in  accurate  terms. 

3.  Where  an  interrogation  point,  (?),  or  a  direct  question  appears 
in  the  directions,  a  corresponding  note  should  appear  in  the  note- 
book.    Oftentimes  the  experiment  itself  does  not  furnish  sufficient 
data  for  answering  some  of  the  questions  asked.     The  necessary  in- 
formation must  then  be  sought  by  referring  to  a  text-book.     To  save 
time  in  such  cases  page  references  to  two  well-known  texts  are  given. 
K.  refers  to  Kahlenberg's  Outlines  of  Chemistry,  and  S.  to  Alexander 
Smith's  General  Chemistry  for  Colleges. 

4.  When  the  word   (Instructions)   appears,  consult  the  instructor 
before  going  further. 

5.  The  chemicals  are  divided  into  three  sets,  each  arranged  alpha- 
betically according  to  the  scientific  names.     The  first  set  contains 
solids  in  bottles  or  jars,  the  second  liquids  in  small  bottles,  and  the 

(7) 


8  Laboratory  Experiments. 

third  liquids  in  large  bottles.  The  bottles  are  arranged  alphabet- 
ically and  their  places  are  numbered  to  facilitate  correct  replacement 
on  the  shelf,  and  particular  care  should  be  taken  to  return  them  to 
their  proper  places.  Read  the  labels  carefully,  and  you  will  not  be 
using  a  sulphate  where  sulphite  is  required,  or  a  concentrated  acid 
where  a  dilute  one  should  be  employed. 

6.     All  students  work  independently,  except  where  the  directions 
say  "Two  students  working  together." 


LAB  ORATORY  EXPERIMENTS  IN  GENERAL  CHEMISTRY. 

CHAPTER  I. 

APPARATUS. 

1.     The  Bunsen  Burner. 

a.  Unscrew  the   chimney  and   examine  the  construction  of  the 
burner.     Make  a  drawing  to  scale  showing  the  various  parts. 

Note  1.  The  purpose  of  drawings  in  the  laboratory  notebook  is 
not  to  represent  a  picture  of  the  apparatus,  but  to  show  its  arrange- 
ment and  operation.  For  this  reason,  and  also  for  simplicity,  sec- 
tional drawings  or  elevations  are  generally  better  than  perspective 
effects.  The  student  should  aim  at  skill  in  making  simple  drawings 
rapidly  and  neatly  with  little  or  no  use  of  a  ruler. 

b.  Attach  the  burner  by  means  of  rubber  tubing  to  the  gas-tap, 
close  the  air-holes  at  the  base,  and  light.     Open  the  air-holes  grad- 
ually and  note  the  effect  on  the  flame.     (Size,   shape,   luminosity, 
steadiness.)     Suggest  a  cause  of  the  difference  in  the  two  flames. 
When  the  air-holes  at  the  base  of  the  burner  are  open  the  gas  should 
burn  with  a  noiseless  blue  flame. 

c.  Explore  the  non-luminous  flame  with  a  platinum  wire  to  locate 
the  relatively  hotter  and  cooler  parts.     (?)     Where  should  an  object 
be  held  to  get  the  greatest  heating  effect?     Hold  a  match  _across  the 
flame  about  1  cm.  above  the  top  of  the  burner;  (?)  now  hold  it  higher 
up.     (?)     Insert   a    match    head    into   the   inner   cone.     Insert   one 
higher  up.     (?)     Show  that  unburnt  gas  exists  in  the  inner  cone  by 
leading  it  out  through  a  glass  tube  and  lighting  it.     What  region  is 
deficient  in  air,  and  which  has  an  excess?     Make  a  scale  drawing  of 
the  non-luminous  flame  showing  its  structure. 

d.  Try  to  vary  the  length  of  the  flames.     Can  you  obtain  a  lumi- 
nous flame  2  cm.  high?     A  non-luminous  flame  of  the  same  height? 
How  could  the  velocity  of  the  gas  in  the  chimney  be  changed  without 
changing  the  pressure  of  the  gas?     How  would  a  change  in  the  veloc- 
ity of  the  chimney  gases  change  the  structure  of  the  flame? 

Note  2.  Perhaps  you  will  find  it  necessary  to  conduct  further 
experiments  before  answering  some  of  the  foregoing  questions.  When 

(9) 


10 


Laboratory  Experiments. 


the  apparatus  for  such  experiments  is  not  in  the  desk  outfit,  it  may 
generally  be  obtained  from  the  store  room  on  temporary  order.  The 
supply  of  apparatus  given  out  on  temporary  order  is  usually  limited. 
It  should  therefore  be  returned  promptly,  so  as  to  make  it  available 
for  other  students.  For  some  forms  of  apparatus  the  store-keeper 
vfiil  require  an  order  from  the  instructor. 

Sometimes  the  instructor  may  have  a  simple  contrivance,  or  he 
may  suggest  a  simple  device  that  may  be  used  in  finding  answers  to 
questions. 

e.  What  volume  of  gas  is  consumed  by  an  ordinary  Bunsen  burner 
per  hour?  What  is  the  approximate  cost  of  running  the  burner  for 
an  hour? 

2.     Capacities  and  Dimensions. 

a.  With  a  triangular  file  make  a  scratch  on  one  of  your  beakers 
just  below  the  flare.     Measure  its  diameter  and  depth  from  mark. 
Calculate   its   capacity   in   cubic   centimeters    (cc.).     Verify    this    as 
under  b. 

b.  With  a  graduated  cylinder  measure  100  cc.  of  water  and  pour 
it  into  one  of  your  beakers.     Readings  on  the  cylinder  are  taken  at 
the  bottom  of  the  meniscus,  and  the  line  of  sight  should  be  in  the 
same  plane.     Next  estimate  the  capacity  of  each  of  your  beakers, 
and  record  your  estimate.     Then  measure  the  volume  of  the  beakers, 
filling  them  to  a  convenient  height  for  use  and  not  to  the  brim.     In 
the  same  way  make  estimates  and  find  the  volumes  of  several  test 
tubes,  the  flasks,  and  bottles  in  your  desk  outfit  of  apparatus.     Re- 
cord the  results  in  tabular  form  somewhat  as  follows: 


Capacities  of  Vessels. 


Name  of  vessel 

Estimated 
volume 

Measured 
volume 

Difference 

Beaker  No.  1  

60  cc. 

72  cc. 

-12  cc. 

Beaker  No.  2 

125  cc. 

115  cc. 

+  10  cc. 

Test  tube  

20  cc. 

Flask,  etc  

Apparatus.  11 

c.  With  a  triangulat  file  make  a  scratch  on  the  neck  of  one  of 
your  flasks.  Fill  it  to  the  mark  with  water  and  measure  its  volume 
with  the  graduated  cylinder.  Now  weigh  the  flask  on  the  platform 
scales.  Again  fill  to  the  mark  with  water  and  weigh.  Find  the 
volume  of  the  flask  by  the  method  of  weighing.  Refer  to  table  4, 
Appendix.  Which  value  for  the  volume  do  you  regard  the  more 
accurate,  and  why?  Express  the  capacity  of  the  flasks  in  ounces, 
d.  Estimate  by  the  eye  the  dimensions  suggested  below  and  re- 
cord your  estimate  and  then  find  the  actual  value  by  measurement. 
Express  results  in  centimeters  (cms.)  or  millimeters  (mms.)  and  frac- 
tional tenths,  and  record  your  results  in  tabulated  form.  Length 
and  diameter  of  test  tubes.  Bore  of  glass  tubing.  Diameter  and 
length  of  rubber  and  cork  stoppers.  Diameter  of  crucible  cover. 
Dimensions  of  laboratory  manual  and  notebook,  etc. 

3.    The  Wash-bottle.     (Figure  1.)     Its  Construction  and  Use. 

a.  Use  a  flask  of  about  500  cc.  capacity,  and  a  two-holed  rubber 
stopper.  Attach  the  nozzle  with  a  short  piece  of  rubber  tubing. 

Working  of  Glass  Tubing.  Cutting.  To  cut  tubing  of  small  size, 
make  a  scratch  at  the  desired  point  with  a  triangular  file.  Hold  the 
tubing  with  both  hands,  place  the  ends  of  the  thumb  nails  together 
opposite  the  scratch,  and  press  forward  with  the  thumbs  and  pull 
back  with  the  hands. 

Bending.  To  bend  glass  tubing,  heat  in  the  ordinary  gas  jet, 
holding  it  lengthwise  in  the  flame,  and  entirely 
in  the  luminous  part,  rotating  it  to  get  an  even 
heating.  Do  not  use  the  Bunsen  flame  unless 
the  burner  is  provided  with  a  wing  top.  When 
the  heated  portion  has  become  soft  enough  to 
bend  by  its  own  weight  take  the  tubing  from 
the  flame  and  bend  it  at  once  to  the  desired 
angle.  While  bending,  sight  along  the  tube  to 
aid  you  in  making  the  two  limbs  come  in  the 
same  plane. 

Fire    Polishing.     The    sharp    edges    of   tubes  Fig.  1. 

should  be  rounded  off  by  heating  the  ends  in  the  Bunsen  flame. 

Drawing.  The  nozzle  is  made  by  heating  a  portion  of  the  tubing 
evenly  in  the  Bunsen  flame  until  it  is  softened,  the  end  portions  being 
held  in  the  hands.  Then  take  the  tube  from  the  flame  and  draw  it 


12  Laboratory  Experiments. 

out  slowly.  When  cool  cut  the  contracted  part  at  the  desired  place 
and  fire  polish  cautiously. 

In  fitting  flasks  with  stoppers  hold  the  flask  firmly  by  the  neck 
and  not  at  the  bottom.  When  inserting  glass  tubing  in  rubber  stop- 
pers or  corks  moisten  the  tubing  or  use  a  little  vaseline,  and  hold  it 
near  the  end  to  be  pushed  into  the  stopper.  Cuts  and  accidents 
too  often  result  from  faulty  methods  of  work. 

No  record  of  this  experiment  is  made  in  the  notebook  up  to  this 
point.  Submit  your  finished  apparatus  to  the  instructor  for  crit- 
icism and  approval.  Finally  fill  the  flask  with  distilled  water  and 
keep  it  in  your  locker  ready  to  be  used  for  washing,  rinsing,  and 
preparing  solutions.  Use  tap  water  for  other  purposes. 

b.  Filtering  and  Washing.  In  the  smallest  flask  or  beaker  place 
a  mixture  of  about  10  g.  of  common  salt  and  2  g.  of  clean  sand,  and 
add  about  30  cc.  of  distilled  water.  Shake  the  mixture  to  aid  so- 
lution. Separate  the  liquid  from  the  undissolved  solid  by  filtering. 
Fold  a  circular  sheet  of  filter  paper  in  halves  and  then  in  quarters. 
Open  the  folded  paper  so  that  a  cone  is  produced,  and  place  the  cone 
in  a  dry  glass  funnel.  Holding  the  paper  in  position  with  one  hand 
blow  a  little  distilled  water  on  it  from  the  wash-bottle  and  with  the 
finger  press  the  paper  firmly  to  the  sides  of  the  funnel.  Next  place 
the  funnel  in  a  filter  stand,  the  stem  extending  into  and  touching 
the  side  of  a  beaker  or  flask  below.  It  is  not  advisable  to  use  the  re- 
tort stand  for  holding  a  funnel.  Why  is  a  wooden  stand  preferred?  To 
avoid  splashing  in  transferring  to  the  filter,  hold  a  glass  rod  against 
the  rim  of  the  beaker  at  the  point  where  the  liquid  is  to  flow  out. 
The  liquid  will  generally  follow  the  rod,  and  by  resting  the  lower 
rounded  end  against  the  filter,  the  liquid  can  be  transferred  without 
loss  or  spattering. 

Rinse  the  beaker  with  a  jet  of  water  from  the  wash-bottle,  and 
when  the  liquid  has  run  out  wash  the  residue  on  the  filter  with  a 
stream  of  water  from  the  wash-bottle  directed  on  the  paper  near 
its  edge.  Continue  the  washing  until  the  taste  of  salt  cannot  be  de- 
tected in  a  drop  of  the  washings.  Suggest  a  more  delicate  test  for 
the  presence  of  salt  in  the  wash  water.  Try  your  test.  (?)  (S.  8.) 

Evaporate  a  portion  of  the  filtrate  on  a  watch-glass,  or  an  evap- 
orating dish,  placed  upon  a  beaker  of  boiling  water.  Use  a  small 
flame  for  heating  and  have  the  wire  gauze  under  the  beaker  close  to 


Apparatus.  13 

the  flame.  Examine  the  residue.  Is  it  still  common  salt?  In 
what  respects  does  it  differ  from  the  original  sample? 

Note  3.  Cleanliness  and  neatness  in  working  are  of  the  utmost 
importance.  Care  should  be  taken  in  little  things.  Attention 
should  be  given  to  the  appearance  of  your  apparatus.  In  setting  up 
apparatus,  ring-stands,  filter  stands,  etc.,  should  be  placed  straight 
on  the  desk.  Iron  rings,  clamps,  etc.,  which  are  not  actually  in  use 
should  be  removed  from  the  retort  stand.  Have  the  rod  of  the  stands 
away  from  you,  not  toward  you.  Keep  your  apparatus  clean  and 
arranged  in  an  orderly  manner  in  your  locker.  If  liquid  has  been 
spilt,  wipe  it  up  directly.  At  the  close  of  a  laboratory  period  clear 
the  desk  and  leave  it  clean  and  dry. 

What  is  the  size  of  the  angle  of  the  filter  funnel? 

4.  The  Simple  Balance.  (Two  students  working  together.) 
The  laboratory  is  provided  with  several  types  of  balances  to  meet 
the  varied  requirements  of  the  experiments.  At  times  the  student 
may  be  uncertain  as  to  which  balance  is  best  suited  for  the  weigh- 
ings to  be  made.  A  moment's  reflection,  however,  will  generally 
result  in  the  proper  choice.  For  example,  when  an  accuracy  in 
weighing  of  a  tenth  per  cent  is  demanded,  the  weight  of  an  object 
of  1000  g.,  when  determined  to  within  1  g.,  meets  the  requirements. 
On  the  other  hand,  the  weight  of  a  1  g.  body  must  be  determined 
to  the  nearest  milligram,  a  thousandth  gram,  to  be  within  the  same 
limit  of  error, — a  tenth  of  one  per  cent.  In  the  former  case  a  bal- 
ance sensitive  to  a  gram  will  answer  the  purpose;  but  for  the  latter 
the  balance  must  be  sensitive  to  a  milligram.  The  platform  scales 
will  safely  carry  a  load  of  2  kilograms  on  each  pan,  and  will  thus 
serve  very  well  for  making  weighings  of  relatively  heavy  forms  of 
apparatus.  The  balance  provided  with  a  pointer  should  be  used 
in  quantitative  experiments,  where  relatively  small  differences  in 
weight  are  to  be  determined  and  the  gross  weight  does  not  exceed 
100  grams.  The  horn-pan  balances  near  the  reagent  shelves  are 
convenient  for  weighing  out  approximate  quantities  of  solid  reagents. 
Look,  before  you  leap:  An  attempt  at  weighing  an  object  of  a  kilo- 
gram mass  on  a  balance  designed  for  a  maximum  load  of  100  g.  may 
ruin  the  balance  and  thus  turn  out  an  expensive  experiment  for  the 
student. 


14  Laboratory  Experiments. 

Directions  for  the  Use  and  Care  of  a  Balance.  (1)  To  release 
the  beam  and  pans  of  the  balance  turn  the  screw  or  lever  on  the  base 
of  the  balance  in  front. 

(2)  To  determine  the  zero  point,  take  half  of  the  total  number 
of  scale  divisions  passed  over  by  the  pointer  in  one  passage;  count 
off  this  number  of  divisions  from  either  end  of  the  swing,  and  use  this 
point  as  the  zero.     The  zero  of  any  one  balance  changes,  and  must 
be  redetermined  every  time  a  weighing  is  made.     The  zero  is  never 
read  by  waiting  for  the  pointer  to  come  to  rest. 

(3)  Handle  all  weights  with  clean  forceps;  never  with  the  fingers. 

(4)  Return  the  weights  to  their  proper  places  in  the  box;  never 
place  them  on  the  desk  or  the  base  of  the  balance. 

(5)  Always  arrest  the  balance  when  not  in  use,  and  every  time 
weights  or  other  objects  are  added  or  removed  from  the  pans.    Never 
take  off  a  weight  whilst  the  beam  is  swinging. 

(6)  How  sensitive  is  the  balance?     For  an  answer,   place  a   10 
gram  weight  on  each  pan  and  note  the  zero  as  before.     Now  increase 
the  weight  on  the  right  hand  pan  by  10  mg.     As  before,  find  the 
position  at  which  the  pointer  would  come  to  rest.     The  difference 
between  the  two  points  of  rest  of  the  pointer  gives  its  deflection  for 
10  mg.,  i.  e.,  the  sensibility  of  the  balance  for  a  10  g.  load.     This 
value  may  be  used  for  determining  weights  less  than  10  mg. 

(7)  Never  weigh  solid  or  liquid  reagents  directly  on  the  balance 
pan.     In  case  of  a  solid  carefully  weigh  a  dry  watch  glass  (or  a  test 
tube)  first,  then  place  the  substance  upon  the  glass  and  weigh  again, 
and  take  the  difference  as  its  weight.     For  a  liquid  a  small  beaker 
(or  a  test  tube)  is  used  instead  of  a  watch  glass. 

(8)  Check  your  count  of  weights  by  the  places  vacant  in  the 
box. 

(9)  Express  weights  in  grams  and  decimals  thereof.     The  weight 
of  an  object  of  17  grams,  8  decigrams,  and  3  centigrams  is  recorded, 
17.83  g. 

(10)  Record  the  weights  in  the  notebook,  or  the  laboratory  man- 
ual, never  on  loose  sheets  of  paper. 

(11)  Be    systematic    in    making    weighings.     Every   change    of   a 
weight  should  bring  you  nearer  to  the  final  result.     Cultivate  the 
estimation  of  weights  and  volumes  in  the  metric  system. 

(12)  Weights  are  generall  placed  on  the  right  hand  pan. 


Apparatus.  15> 

Practice. 

(1)  Ascertain  the  weight  of  several  objects  designated  by  the  in- 
structor.    Weigh  each  separately  and  then  find  the  combined  weight 
by  weighing.     If  the  sum  of  the  weights  of  the  individual  objects 
differs  from  the  total  weight  found  by  more  than  0.02  g.  repeat  the 
weighings. 

(2)  Weigh  a  five-cent  piece. 

(3)  Set  up  a  burette  as  shown  in  Fig.  33,  K.     Clean  the  burette. 
(Instructions).     A  clean  burette  will  not  leave  streaks  of  water  or 
drops  on  its  walls  when  liquid  is  withdrawn.     Fit  it  with  a  short 
piece  of  rubber  tubing  and  glass  nozzle.     The  flow  of  liquid  is  con- 
trolled by  a  pinch-cock  or  by  placing  a  glass  bead  in  the  middle  of 
the  rubber  tube.     When  using  the  burette  make  sure  that  the  rub- 
ber tube  and  nozzle  are  free  from  air-bubbles,  and  there  is  no  leak 
at  the  nozzle. 

Weigh  a  small  dry  beaker.  From  a  burette  measure  into  the  beak- 
er about  20  cc.  of  distilled  water.  Make  the  readings  on  the  burette 
carefully,  observing  the  lower  line  of  the  meniscus,  and  estimating 
tenths  of  a  division.  Allow  about  a  minute  between  burette  read- 
ings for  the  water  to  drain  down  the  wall.  Note  carefully  the  vol- 
ume of  water  measured  into  the  beaker,  expressing  it  in  cubic  cen- 
timeters, e.  g.,  17.85  cc.  Now  weigh  the  beaker  and  contents.  From, 
the  data  obtained  compute  the  weight  of  1  cc.  of  water.  Criticise 
the  result. 

Make  a  systematic  record  of  the  data: 

Weight  of  beaker  and  water =33.47     g. 

Weight  of  beaker  empty =  15.74     g. 

Weight  of  water =17.73     g. 

Volume  of  water =17.95   cc~ 

Weight  of  1  cc.  of  water =  0.988  g. 

Allow  the  beaker  to  stand  uncovered  for  30  minutes,  then  weigh 
again.  Compute  the  percentage  loss  in  weight.  Mention  three 
factors  that  influence  the  rate  of  evaporation  of  water  from  an  open. 
vessel. 

5.  Identification  of  Substances.  (Two  students  working  to- 
gether.) 

Bring  specimens  of  charcoal,  saltpetre  (potassium  nitrate),  and, 
roll  sulphur  on  watch  glasses  from  the  side-shelf. 


16  Laboratory  Experiments. 


a.  Examine  the  sulphur  carefully,  noting  among  its  specific  prop- 
erties,  color,  odor,   hardness,   and  taste.     In  testing  the  taste  of  a 
substance  touch  only  a  minute  quanity  to  the  tongue.       Is  sulphur 
soluble  in  water?     Test  its  solubility  in  carbon  disulphide,  using  a 
dry  test  tube.     (Caution: — Carbon   disulphide  is  very  inflammable. 
Keep   it  away  from  the  flame.)     The  relative  solubility  of  sulphur 
in  water  and  carbon  disulphide  can  be  judged  by  the  eye.     To  ascer- 
tain whether  sulphur  is  appreciably  soluble  in  water  filter  the  mix- 
ture after  shaking  well,  catch  a  few  drops  on  a  watch  glass,  and  evap- 
orate on  a  water  bath.     If  the  mixture  settles  readily  pour  off  a  few 
drops  of  the  clear  liquid  directly  upon  a  watch  glass  and  evaporate. 
Is  the  stain  upon  the  watch  glass  any  greater  than  the  same  quan- 
tity of  solvent  leaves?     Heat  a  small  piece  of  sulphur  in  a  test  tube. 
(?)     Does  it  appear  to  melt  at  a  lower  temperature  than  potassium 
nitrate?      Compare  its  melting  point  with  that  of  potassium  nitrate. 
(K.    178,  349;  S.    249,    366.)     Place    a    small  piece  of  sulphur  in  a 
deflagrating  spoon  and  heat  it  in  a  Bunsen  flame.     (?)      Does  it  take 
fire  before  it  melts?     When  it  begins  to  burn  remove  the  spoon  from 
the    flame  and  with  the  hand  waft  some  of  the    gas   formed  toward 
the  nose  to  test  its  odor. 

Examine  the  specimens  of  charcoal  and  saltpetre,  making  sim- 
ilar tests  and  finally  tabulate  the  properties  of  the  three  substances. 
Include  in  this  table  some  other  characteristic  properties  of  these 
substances,  such  as  the  melting  points,  boiling  points,  densities, 
specific  heats,  electrical  conductivity.  (K.  214,  348,  178;  S.  318, 
365,  248.) 

b.  Gunpowder   is    made    from    charcoal,    saltpetre,    and    sulphur. 
place  about  2  g.  of  the  ordinary  powder  in  a  test  tube,  add  10  cc. 
of  water,  shake  well,  and  also  warm  gently.     Filter;  wash  the  res- 
idue on  the  paper  with  a  jet  of  water  blown  from  the  wash-bottle. 
Evaporate  the  filtrat  on  a  water  bath.       Examine    the  residue  and 
identify  it.     Dry  the  filter  paper  and  its  black  residue  by  spreading 
it  out  on  the  sand  bath  and  warming  gently.     Transfer  the  residue 
to  a  dry  test  tube  and  shake  it  with  about  3  cc.  of  carbon  disulphide. 
Filter  and  allow  the  filtrate  to  evaporate  without  heating  it.     De- 
scribe and  name  the  residue.     Suggest  a  reason  for  drying  the  res- 
idue before  shaking  it  with  carbon  disulphide.     Examine  the  black 
residue  on  the  paper  and  compare  its  properties  with  those  of  finely 


Apparatus.  17 

powdered  charcoal,  obtained  by  grinding  a  piece  of  charcoal  in  the 
mortar.  What  is  the  probable  cause  for  any  difference  in  the  be- 
havior of  the  residue  and  the  powdered  charcoal? 

Does  any  chemical  change  occur  during  the  manufacture  of  gun- 
powder? 

Estimate  the  relative  quantities  of  the  components  present  in  the 
specimen  of  gunpowder  studied,  and  make  comparisons  with  the 
composition  of  the  ordinary  powder.  (K.  349;  S.  366.) 


Chem.— 2 


CHAPTER  H. 

HYDROGEN. 

Note  4.  Before  commencing  an  experiment  read  carefully  the 
directions  to  grasp  the  purpose  of  the  experiment,  and  to  plan  the 
work  involved.  Record  in  your  notebook  what  you  observe  and 
guard  against  entering  notes  about  what  you  think  you  should  ob- 
serve. If  you  find  that  your  observations  do  not  seem  to  be  in  ac- 
cord with  the  text-book  or  the  statements  of  the  lectures,  read  the 
directions  again,  reflect  a  moment,  and  then  try  to  clear  up  the  dis- 
crepancies that  appear  to  exist  by  repeating  the  experiment,  and 
finally  consulting  the  instructor.  In  this  way,  the  scientific  method 
of  dealing  with  the  problem  in  hand  will  be  acquired,  and  the  final 
solution  of  the  difficulty  with  its  thrill  of  satisfaction  may  prove  of 
more  educational  value  than  several  experiments  without  compli- 
cations. 

6.    Interaction  of  Metals  and  Acids. 

a.  Set  up  about  a  dozen  test  tubes  in  the  test  tube  rack.  Place 
a  few  pieces  of  each  of  the  following  metals  in  separate  tubes: — Alu- 
minium (turnings),  copper  (chips),  iron  (filings),  lead  (clippings), 
magnesium  (wire  or  ribbon),  platinum  (wire),  tin  (granulated),  zinc 
(gran.),  zinc  (dust).  Pour  into  the  graduated  cylinder  25  cc.  of  con- 
centrated hydrochloric  acid  and  dilute  it  with  an  equal  volume  of 
water,  stirring  the  mixture  to  insure  thorough  mixing.  Add  5  cc.  of 
the  diluted  acid  to  one  of  the  metals,  and  watch  the  action  carefully, 
and  record  your  observations.  If  little  or  no  action  occurs  in  the 
cold,  heat  in  a  small  Bunsen  flame.  (?)  Follow  the  same  method 
of  experimenting  with  the  other  metals.  When  gas  arises  from  the 
mixture,  note  where  it  first  appears.  Test  the  issuing  gas  by  bring- 
ing a  light  to  the  mouth  of  the  test  tube.  If  a  pop  follows  or  a  flame 
is  seen,  the  gas  is  hydrogen.  When  the  interaction  is  slight  the  gas 
set  free  may  not  test  for  hydrogen  by  this  method.  The  formation 
of  small  bubbles  of  gas  on  the  metal  after  heating  indicates  hydro- 
gen. In  such  cases  a  sample  of  the  gas  may  be  collected  in  a  small 

(18) 


Hydrogen. 


19 


test  tube  filled  with  water,  and 
inverted  over  water,  as  shown  in 
Fig.  2.  The  whole  apparatus 
may  be  conveniently  mounted 
on  a  ringstand.  If  the  gas  sam- 
ple thus  collected  appears  to  be 
hydrogen  give  a  reason  for  not 
detecting  it  as  it  escaped  from 
the  open  tube. 

Arrange  the  metals  in  the  or- 
der of  activity.  In  what  respects 
does  your  order  differ  from  the 
electrochemical  series?  (K.  435;  Fig.  2. 

S.  245.)  Suggest  a  reason  for  some  of  the  differences.  Where  would 
you  place  hydrogen  in  this  series?  What  probably  determined  the 
order  on  which  the  alchemic  symbols  of  Fig.  3  are  arranged? 


10 


Au 


Pt 


Ag 


Hg          Cu 
Fig.  3. 


Pb 


ii 
^ 


Fe 


Note  5.  Throw  matches,  wet  filter  papers  and  other  solid  ob- 
jects in  the  waste  jars  under  the  desk.  Do  not  throw  them  into  the 
sink. 

What  percentage  (approximate)  of  the  dilute  acid  used  is  water? 
(K.  50.)  Is  all  of  the  hydrogen  of  the  acid  displaced  in  the  course 
of  the  interaction  with  an  excess  of  zinc?  If  not,  formulate  your 
ideas  about  the  proportion  of  the  total  hydrogen  that  may  be  dis- 
placed from  dilute  acid  by  zinc. 

(Optional.)  Assuming  that  the  diluted  acid  used  contains  20  per 
cent  of  hydrogen  chloride,  compute  the  percentage  of  hydrogen  in 
the  solution,  taking  the  composition  of  water  as  11.14  per  cent  hydro- 
gen, and  88.86  per  cent  oxygen,  and  that  of  hydrogen  chloride  as 
2.77  per  cent  hydrogen  and  97.23  chloride.  In  these  experiments 
what  percentage  of  the  total  hydrogen  of  the  solution  can  be  dis- 
placed by  a  metal  like  zinc? 


20  Laboratory  Experiments. 


b.  Prepare  some  dilute  sulphuric  acid  by  pouring  5  cc.  of  the  con- 
centrated acid  into  15  cc.  of  water.     Try  its  interaction  with  small 

pieces  of  zinc  and  copper,  in  separate  test  tubes.  (?) 
Just  cover  a  few  pieces  of  zinc  in  a  test  tube  with  con- 
centrated sulphuric  acid.  (?)  Make  a  paper  holder 
for  the  test  tube  and  heat  the  mixture.  (?)  Test  the 
odor  of  the  products.  (Exp.  5a.)  What  products  of 
the  reaction  can  you  identify?  Describe  one  of  the 
products  that  is  new  to  you.  Try  also  the  action  of 
Fig.  4.  dilute  and  concentrated  sulphuric  acid  on  pieces  of 
copper. 

c.  Try  the  action  of  nitric  acid,  concentrated  and  dilute,  on  sev- 
eral of  the  metals,  zinc  or  copper,  and  iron  (wire),  tin.     (?)     Is  hydro- 
gen obtained?     Exercise  care  in  working  with  nitric  acid.     It  leaves 
yellow  stains  on  the  hands. 

d.  Try  the  action  of  acetic  acid  on  zinc  and  magnesium.     (?) 

e.  Fuse  the  end  of  your  platinum  wire  into  a  short  piece  of  glass 
rod.     Wrap  the  wire  around  a  piece  of  zinc,  and  dip  it  into  very 
dilute  sulphuric  acid  (1:20).     Do  any  gas  bubbles  appear  on  the  wire? 
Is  it  dissolving?     What  part  does  the  platinum  wire  take  in  this  re- 
action?    (S.  66.)     Does  gas  form  on  the  wire  (a)  when  the  zinc  and 
platinum  do  not  touch  each  other?  (b)  when  the  ends  outside  the 
acid  are  in  contact? 

What  is  the  commercial  name  of  hydrochloric  acid?  Of  sulphuric 
acid? 

Modify  the  experiment  by  using  copper,  carbon,  or  gold  instead  of 
the  platinum  wire.  (?) 

Note  6.  When  acid  gets  upon  the  clothing,  apply  ammonium  hy- 
droxide at  once.  Burns  should  be  rubbed  with  a  paste  of  sodium 
hydrogen  carbonate  (baking  soda)  and  water.  It  is  well  to  dress 
painful  burns  with  a  preparation  of  lime  water  and  sweet  oil.  Cuts 
should  be  washed  at  once  with  running  water,  and  dressed  with  boric 
acid,  or  a  solution  of  lysol.  Obtain  the  assistance  of  an  instructor 
in  case  of  injury. 


Hydrogen. 


21 


Fig.  5. 


7.     Preparation  and  Properties  of  Hydrogen. 

a.  Fit  a  250  cc.  flask  with  a  thistle  tube  and 
a  delivery  tube.     Put   30-50  g.   of  granulated 
zinc  in  the  flask  and  pour  50  cc.  of  dilute  hy 
drochloric  acid  (1:4)    on  it.  Collect  some  of  the 
gas  issuing  from  the  delivery  tube  over   water 
in  the  pneumatic  trough  (K.  16;  S.  48.)     Test 
the     gas    to    see    whether    it    is  free  from  air 
(source?)  by  bringing  a  light    near  the  mouth 
of  the  test  tube.     Continue  the  tests  with  fresh 
samples  until  the   gas  in   the  test  tube  burns 
quietly. 

Fill  one  of  the  wide-mouth  bottles  with  hy- 
drogen, and  holding  it  mouth  downward,  thrust 
into   the   flask   a   burning   splinter.     Does   the   splinter   continue   to 
burn?     Explain. 

Show  by  an  experiment  that  hydrogen  is  lighter  than  air.     (?) 

Cover  a  wide-mouth  bottle  filled  with  hydrogen  with  a  sheet  of 
dry  filter  paper,  and  place  it  mouth  downward  on  a  piece  of  gauze 
supported  on  a  ring  of  the  retort-stand,  and  leave  it  for  two  minutes. 
Then  test  for  hydrogen.  (?)  Try  the  experiment  substituting  a  wet 
sheet  of  paper.  (?)  Try  it  also  with  a  sheet  of  note-paper.  (?) 
Explain  the  results. 

Replace  the  delivery  tube  with  a  glass  nozzle.  If  the  action  in  the 
generator  has  slackened  pour  5  cc.  of  concentrated  hydrochloric  acid 
through  the  thistle  tube.  Hold  a  cold,  dry  beaker  against  the  nozzle. 
If  moisture  is  deposited  what  is  its  probable  source?  Connect  a 
U-tube  filled  with  fused  calcium  chloride  between  the  nozzle  and  the 
outlet  tube.  Hold  the  cold  beaker  against  the  jet  once  more.  (?) 
When  the  issuing  gas  is  free  from  moisture  test  it  to  see  that  it  is  not 
explosive.  Then  light  the  jet  and  again  hold  a  dry  cold  beaker 
near  the  flame.  (?)  Why  is  the  hydrogen  dried  for  this  experiment? 

Burn  dried  illuminating  gas  at  the  nozzle,  and  hold  a  dry,  cold 
beaker  over  the  flame  which  should  be  small.  (?)  Explain. 

Draw  from  memory  a  figure  of  the  Kipp  apparatus  used  in  the 
class-room  for  preparing  hydrogen. 

b.  (Two  students  working  together.)     Fit  a  hard  glass  tube  at 
least  20  cm.  long  with  corks  and  tubes  as  shown  in  the  accompanying 


22 


Laboratory  Experiments. 


figure.  Dry  specimens  of  ferric  ox- 
ide, silicon  dioxide  (sand),  and  cu- 
pric  oxide  by  heating  in  a  crucible. 
Put  one  of  the  oxides  in  a  porcelain 
boat  and  place  it  in  the  tube.  Now 
connect  the  tube  to  a  source  of  dry 
hydrogen  and  pass  a  stream  of  the 
gas  through  the  tube.  Test  the  is- 
suing gas  to  see  that  it  is  free  from 
air.  Then  heat  the  tube,  cautiously 
at  first  by  moving  the  flame  back 
Fig.  6.  and  forth,  and  later,  strongly  until 

the  boat  is  red  hot.     What  evidence  does  the  experiment  furnish  that 
a  chemical  change  is  in  progress  during  the  heating? 

Conduct  similar  experiments  with  the  other  oxides.  Name  the 
new  products  formed. 

What  property  of  hydrogen  do  these  experiments  illustrate?  As- 
sign a  title  to  this  paragraph  in  your  notebook. 

c.  (Optional;  Two  students  working  together.)  Perform  Exp. 
13  d.  Transfer  the  weighed  residue  of  anhydrous  cupric  sulphate 
to  the  porcelain  boat.  Guard  against  loss  of  material.  Place  the 
boat  in  the  reduction  tube  and  pass  dried  hydrogen  over  it  till  con- 
stant weight  is  obtained.  Identify  the  residue.  What  does  the  loss 
in  weight  represent?  During  the  reduction  note  the  odor  of  the  is- 
suing gas.  (?)  What  percentage  of  blue  vitriol  remains  as  residue? 


CHAPTER  III. 

OXYGEN. 

8.     Sources. 

Oxygen  is  one  of  the  constituents  of  the  following  substances  and 
the  percentages  show  the  approximate  quantity  present  in  each: 

Lead  dioxide,  13  per  cent. 

Potassium  nitrate,  47  per  cent. 

Silicon  dioxide  (sand),  53  per  cent. 

Barium  peroxide,  19  per  cent. 

Sugar,  51  per  cent. 

Manganese  dioxide,  36  per  cent. 

Calcium  carbonate  (marble  or  limestone),  48  per  cent. 

Heat  2-3  g.  of  each  of  these  substances  in  a  dry,  hard  glass  test 
tube.  Observe  whether  gas  is  given  off  and  in  each  case  test  it  by 
inserting  into  the  test  tube  a  glowing  splinter.  Try  the  blast  lamp 
if  the  Bunsen  flame  fails  to  give  results.  If  drops  of  liquid  collect 
on  the  sides  of  the  test  tube  during  heating  hold  it  in  a  nearly  hori- 
zontal position.  Why?  Make  a  record  of  the  changes  noted  in  the 
course  of  the  heating.  Describe  the  residues.  Arrange  the  sub- 
stances studied  in  the  order  of  their  stability.  (S.  81.) 

Note  7.  To  dry  test  tubes  quickly  warm  them  in  the  flame,  and 
then  blow  or  draw  a  stream  of  air  into  the  tube  through  a  glass  tube 
reaching  nearly  to  the  bottom. 

To  repair  a  test  tube  seal  on  a  short  glass  rod  in  the  blast-lamp. 
(Instructions.)  Now  heat  the  tube  evenly  a  little  above  this  joint. 
When  the  glass  has  softened  pull  off  the  end  part.  Soften  the  sealed 
end  of  the  tube  again,  remove  it  from  the  flame,  and  holding  it  ver- 
tically, blow  to  round  out  the  end.  Repeat  the  softening  and  blow- 
ing until  the  wall  is  of  nearly  uniform  thickness.  Some  persons  pre- 
fer to  heat  the  broken  end  of  the  tube  directly  in  the  flame  until  it 
closes  up.  The  lump  of  glass  in  the  bottom  is  then  blown  out  as 
just  stated.  Test  tubes  of  soft  glass  are  readily  repaired.  Hard  glass 
is  more  difficult  to  work. 

(23) 


24 


Laboratory  Experiments. 


9.    Preparation  and  Properties. 

a.  Mix  on  paper  about  5  g.  of  potassium  chlorate  and  3  g.  of  man- 
ganese dioxide.  Place  the  mixture  in  a  test  tube  fitted  with  a  one- 
hole  cork  and  delivery  tube.  Test  the  apparatus  to  see  that  it  is 
air-tight.  Then  arrange  the  parts  as  shown  in  Fig.  7.  Evolve  the 
oxygen  moderately  slowly  by  regulating  the  heating.  Collect  the 


Fig.  7. 

gas  over  water,  filling  five  bottles.  Keep  the  residue  in  the  test  tube 
and  treat  it  as  directed  under  c. 

b.     Test  the  odor  of  the  gas.     Inhale  some  of  it.     (?) 

Place  a  small  piece  of  sulphur  in  a  deflagrating  spoon,  and  burn  it 
in  one  of  the  bottles  of  gas.  (?)  Identify  the  product  by  its  odor. 
(Exp.  5a.) 

Pour  a  little  water  into  the  bottle,  close  the  mouth  with  the  hand 
and  shake.  (Object  of  this?)  Test  the  liquid  with  litmus  solution 
or  blue  litmus  paper.  (?) 

What  is  an  indicator?     (K.  130;  S.  242.) 

Burn  a  little  red  phosphorus  in  another  bottle,  and  make  the  lit- 
mus test.  (If  yellow  phosphorus  is  used,  GREAT  CARE  must  be 
taken  not  to  get  it  on  the  hands.  It  catches  fire  easily,  and  causes 
severe  burns.  Cut  it  under  water  and  handle  with  forceps.) 

Lower  a  piece  of  glowing  charcoal  into  a  bottle  of  the  gas.  (?)  When 
burning  has  ceased,  add  a  little  distilled  water  (why  not  tap  water?), 


Oxygen. 


25 


cover  the  bottle  with  a  glass  plate  and  shake.  Add  some  clear  lime- 
water,  and  shake  again.  Interpret  the  lime-water  test.  What  is  the 
chemical  name  of  lime-water? 

Repeat  the  burning  of  charcoal  and  make  the  litmus  test.     (?) 

c.  Add  distilled  water  to  the  black  residue  in  the  test  tube,  warm 
gently,  and  transfer  to  a  filter.  Test  the  filtrate  by  adding  a  few 
drops  of  silver  nitrate  solution.  (?)  Similarly,  add  a  few  drops  of 
silver  nitrate  solution  to  a  solution  of  potassium  chloride.  (?)  To  a 
solution  of  potassium  chlorate  add  a  few  drops  of  silver  nitrate.  (?) 
What  inference  may  be  drawn  from  these  tests? 

Wash  the  residue  on  the  filter  until  a  few  drops  of  the  washings  pro- 
duce little  or  no  turbidity  with  a  few  drops  of  silver  nitrate  solution. 
Now  dry  the  residue.  Heat  it  in  hard  glass  test  tube  in  the  blast- 
lamp  and  try  the  glowing  splinter  test.  (?)  If  oxygen  appears  what 
is  its  probable  source? 

10.     Weight  of  a  Liter  of  Oxygen.     (Quant.)*     (Two  students). 

In  this  experiment  oxygen  is  evolved  by  heating  potassium  chlorate. 
Powder  some  of  the  cholrate  and  dry  it  on  a  watch  glass  by  gentle 
heating  on  the  sand  bath.  Weigh  a  hard  glass  test  tube  and  place  in 
it  about  1.5  g.  of  the  dried  chlorate.  Weigh  the  tube  and  chlorate 
together.  Fit  up  the  apparatus  shown  in  Fig.  8,  using  the  bottle  of 


Fig.  8. 


*  The  first  paragraph  of  Expt.  4,  p.  13,  will  guide  you  in  selecting 
the  proper  balance  for  the  different  weighing  operations  in  this  experi- 
ment. 


26  Laboratory  Experiments. 


about  one  liter  capacity  for  the  aspirator.  All  joints  must  be  per- 
fectly air-tight.  This  can  be  ascertained  by  having  the  delivery  tube 
leading  to  the  beaker  full  of  water  and  then  opening  the  clip.  If  the 
apparatus  is  air-tight  only  a  few  drops  of  water  will  flow  out  at  first. 
If  it  is  not  air-tight  water  will  continue  to  trickle  out.  At  the  begin- 
ning of  the  experiment  have  the  exit  tube  full  of  water  and  the  aspi- 
rator at  least  two-thirds  full.  Allow  the  exit  tube  to  drop  to  the  bot- 
tom of  the  beaker,  which  should  contain  a  little  water  and  be  of  from 
400-500  cc.  capacity.  Open  the  clip  and  by  raising  the  beaker  equalize 
the  levels  of  the  water  in  the  aspirator  and  the  beaker.  (Why?)  Close 
the  clip;  empty  the  beaker  and  then  replace  it.  After  opening  the  clip 
again,  heat  the  test  tube  slowly,  gradually  raising  the  temperature. 
In  the  course  of  the  decomposition  of  the  chlorate  smoke  is  oftentimes 
seen.  This  consists  of  solid  particles  and  its  loss  from  the  tube  should 
be  guarded  against.  (Why?)  If  the  test  tube  is  of  a  rather  small 
bore  be  watchful  lest  the  passage  for  the  gas  becomes  stopped.  By 
heating  near  the  upper  level  of  the  salt  a  free  passage  for  the  gas  gen- 
erally remains.  When  300-500  cc.  of  the  water  has  been  driven  out 
of  the  aspirator  into  the  beaker,  stop  heating,  allow  to  cool  to  the  tem- 
perature of  the  room,  equalize  the  levels  and  close  the  clip.  Determine 
the  volume  of  the  water.  (See  Exp.  2c.)  Weigh  the  test  tube  with 
the  residue  in  it.  Read  the  barometer,  the  temperature  near  the 
barometer,  and  the  temperature  of  the  water.  You  now  have  the 
weight  of  a  known  volume  of  oxygen  standing  over  water  at  the  tem- 
perature and  barometric  pressure  observed.  Reduce  the  volume  to 
0°  and  760  mm.  pressure,  allowing  on  the  corrected  barometric  pres- 
sure for  the  tension  of  water  vapor.  (Appendix,  tables  5,  6  and  7.) 
Calculate  the  weight  of  a  liter  of  oxygen  at  standard  temperature  and 
pressure. 

What  is  the  volume  occupied  by  32  g.  of  oxygen? 

Should  account  be  taken  of  the  volume  of  the  air  originally  present 
in  the  aspirator? 

Is  the  apparatus  of  this  experiment  suitable  for  determining  the 
weight  of  a  liter  of  hydrogen?  What  changes  in  the  apparatus  would 
you  introduce  and  why? 


Oxygen.  27 


Data. 

Wt.  of  tube  with  chlorate g. 

Wt.  of  tube  and  residue 

Wt.  of  oxygen g. 

Wt.  of  beaker  with  water g. 

Wt.  of  beaker 

Wt.  of  water g. 

Temp,  of  water 

Temp,  near  barometer 

Barometric  reading mm. 

Correction 

Barometer  (corr.) 

Tension  of  water  vapor 


Partial  pressure  of  oxygen mm. 

Volume  of  oxygen,  S.  T.  P cc. 

Weight  of  1  liter  of  oxygen g. 

Volume  of  32  g.  of  oxygen,  S.  T.  P cc. 

11.  Melt  a  small  quantity  of  potassium  chlorate  cautiously  in  a 
test  tube  clamped  in  a  vertical  position.  Use  a  very  small  Bunsen 
flame  for  heating.  Can  the  chlorate  be  melted  without  decomposing? 
Throw  a  pinch  of  manganese  dioxide  into  the  molten  chlorate.  (?) 
Apply  the  glowing  splinter  test.  Interpret  the  result.  (S.  54.)  As- 
sign a  name  to  this  experiment  in  your  record. 


CHAPTER   IV. 


WATER. 

12.  Some  Properties  of  Water. 

a.  Purity.    Taste  some  distilled  water.  Could  you  distinguish  it  from 
tap  water  by  this  property?  Suggest  two  other  methods  of  distinguishing 
it  from  tap  water.     Try  your  methods.     Devise  a  method  of  testing 
distilled  water  for  the  presence  of  dissloved  air,  and  suggest  a  method 
of  removing  the  air  dissolved  in  it.     Compare  the  tints  produced  on 
litmus  paper  respectively  by  distilled  and  tap  water. 

b.  Union  with  Oxides.     In  separate  test  tubes  place  small  quanti- 
ties of  the  following  oxides:     Cupric  oxide,  quick-lime,  ferric  oxide, 
phosphorus  pentoxide,  barium  oxide  or  peroxide.     Add  a  few  cc.  of 
distilled  water  to  each  and  shake.     Filter  and  test  the  filtrates  with 
litmus  solution  or  litmus  paper.     (?) 

What  other  acid  forming  oxides  have  been  examined  in  earlier 
experiments? 

What  class  of  elements  furnish  acid  forming  oxides?     (K.  125;  S.  81.) 

13.  Hydrates. 

a.  (Two  students.)  Gently  heat  specimens  of  barium  chloride, 
sodium  chloride,  Glauber's  salt,  potassium  nitrate,  alum,  potassium 
dichromate,  and  sugar.  Note  the  results  that  occur  during  the  course 
of  the  heating  and  record  the  results  in  tabular  form,  under  the  follow- 
ing headings: 


Amount  of  de- 

Substance 
heated. 

Changes   during 
heating. 

posit  in  cool 
portion  of 
tube. 

Appearance 
of  residue. 

Are  all  crystalline  substance  hydrates?  Which  of  the  substances 
examined  would  you  classify  as  hydrates?  Does  sugar  lack  any  of  the 
characteristics  of  hydrates?  (S.  82;  K.  257.) 

(28) 


Water.  29 


b.  (Two  students  working  together.)     Place  several  small  crystals 
of  blue  vitriol  in  a  test  tube,  cover  them  with  concentrated  sulphuric 
acid,  and  note  the  changes  that  occur  in  the  course  of  a  half  hour. 
Meanwhile  heat  several  crystals  of  blue  vitriol  in  an  evaporating  dish 
or  a  test  tube.     Set  aside  a  small  portion  of  the  resulting  white  powder 
on  a  watch  glass.     (?)     Boil  the  remainder  with  sufficient  water  to 
dissolve  it.     Avoid  an  excess  of  water.     Set  the  solution  aside  to  cool 
(?)     Drain  off  the  sulphuric  acid  from  the  crystals  and  add  a  few  drops 
of  water  to  the  residue  in  the  test  tube.     (?). 

c.  (Optional;  Quant.)     Weigh  out  accurately,  about  2  g.  of  gypsum 
in  a  porcelain  crucible      Place  the  crucible  on  a  clay  triangle,  and  heat 
to  redness.     When  cool,  weigh,  and  then  heat  a  second  time  and  weigh. 
When  the  weight  is  constant  calculate  the  percentage  loss  in  weight. 
Record  the  data  in  systematic  form  as  suggested  by  the  following 
outline: 

Hydrate :    Gypsum. 

Wt.  of  crucible  with  gypsum.    =   9.485    g. 
Wt.  of  crucible  =   7.620    g. 


Wt.  of  salt  taken  =   1.865    g. 

Wt.  of  crucible  with  residue  =  g.          (First  weighing) 

Wt.  of  crucible  with  residue  =  (Second  weighing) 

Wt.  of  crucible  with  residue  =  (Third  weighing) 

Wt.  of  crucible 

Wt.  of  residue  =  g. 

Loss  in  weight  =  g. 

Percentage  loss  in  weight per  cent. 

What  does  the  loss  in  weight  represent? 

d.  (Quant.;  optional.  Two  students.)  Weigh  a  crucible  with  cover. 
Place  in  the  crucible  about  2  g.  of  blue  vitriol.  Crystals  having  a  white 
crust  should  be  rejected.  Cover  and  weigh.  Nowjjupport  the  covered 
crucible  over  a  Bunsen  flame  so  that  the  top  of  the  flame  comes  within 
3  cm.  of  the  bottom  of  the  crucible.  Heat  thus  for  ten  minutes.  Allow 


30  Laboratory  Experiments. 

to  cool;  and  then  weigh.  Heat  for  five  minutes  longer;  weigh  again 
when  cool.  When  no  further  change  in  weight  occurs,  use  the  data 
to  calcualte  the  percentage  loss  in  weight.  What  does  the  loss  in 
weight  represent? 

Suggest  a  reason  for  not  heating  the  blue  vitriol  as  highly  as  the 
gypsum  in  these  experiments. 

14.     Solubility.     (Quant.) 

Pulverize  about  15  g.  of  potassium  dichromate.  Prepare  a  saturated 
solution  of  the  substance  by  shaking  (Patience!)  it  in  a  stoppered  flask 
with  50  cc.  of  water.  Take  the  final  temperature.  Now  decant  the 
clear  solution  into  a  burette,  and  measure  a  definite  volume,  20-30  cc. 
into  a  weighted  exaporating  dish,  and  weigh  again.  Evaporate  the 
solution  on  a  water  bath,  completing  the  final  drying  by  gentle  warm- 
ing over  a  free  flame.  (Instructions.)  Weigh  the  residue,  and  from 
the  data  in  hand  compute  the  quantity  of  the  salt  that  would  be  dis- 
solved by  100  g.  of  water  at  the  observed  temperature.  Calculate  also 
the  quantity  of  salt  that  a  liter  of  this  solution  would  contain.  Make 
a  tabulated  statement  of  your  data,  and  record  the  solubility  deter- 
minations of  three  other  students.  Compare  the  results  with  a  view 
of  accounting  for  some  of  the  differences  in  the  values  as  compared 
with  your  result.  What  was  the  density  of  your  solution? 

The  solubility  of  potassium  dichromate  in  water  at  certain  tempera- 
tures is  given  in  the  following  table: 

Temperature.  Grams  of  K,Cr2O7 

in  100  g.  of  water. 

0°  5.0 

10  8.5 

20  13.1 

40  29.2 

60  50.5 

Represent  the  above  data  by  a  graph,  plotting  temperatures  on  the 
axis  of  abscissae,  and  the  corresponding  solubilities  as  ordinates. 
Draw  the  solubility  curve.  (K.  413;  S.  104).  Represent  upon  the 
diagram  your  solubility  determination  and  the  results  of  three  other 
students. 


CHAPTER  V. 


EQUIVALENT    WEIGHTS,    FORMULAS,    EQUATIONS. 

15.     Composition  of  Magnesium  Oxide.     (Quant.) 

Weigh  a  crucible  with  cover.  Clean  a  piece  of  magnesium  ribbon 
about  a  meter  long  with  a  bit  of  emery  cloth.  If  the  magnesium  is 
bright  it  may  be  used  without  cleaning.  Roll  the  ribbon  up  loosely 
but  small  enough  to  lie  in  the  crucible.  Place  it  in  the  crucible,  cover 
and  reweigh.  Place  the  crucible  and  contents  on  a  triangle  and  heat 
it  at  some  point  where  the  magnesium  touches  it.  As  soon  as  the 
magnesium  begins  to  glow  regulate  the  supply  of  air  by  sliding  the 
cover  on  and  off.  Allow  as  much  air  to  enter  as  is  possible  without 
letting  the  white  oxide  escape.  A  little  loss  will  undoubtedly  occur, 
but  great  care  should  be  taken  to  have  the  loss  small.  After  the  burn- 
ing has  nearly  ceased,  remove  the  cover  entirely,  but  guard  against 
the  loss  of  oxide  that  clings  to  it.  Now  heat  the  crucible  to  bright 
redness  for  five  minutes  in  the  top  of  the  flame.  Toward  the  end  of  the 
heating  the  solid  may  be  carefully  stirred  with  a  pointed  glass  rod. 
It  is  well  to  heat  the  rod  to  keep  it  from  breaking  when  it  touches  the 
hot  oxide.  What  is  the  necessity  of  stirring  the  powder?  Allow  the 
crucible  to  cool  and  then  weigh.  Continue  the  heating,  cooling  and 
weighing  until  a  constant  weight  is  obtained.  The  contents  of  the 
crucible  should  finally  be  perfectly  white.  Interpret  the  gain  in  weight. 

From  your  data  calculate  the  weight  of  magnesium  combining  with 
8  parts  of  oxygen. 

Wt.  of  Ox.  found     :     Wt.  of  Mg.  taken     :     8     :     x. 

This  gives  x,  the  equivalent  weight  of  magnesium. 

Record  the  data  obtained  in  the  experiment  in  tabular  form  as 
follows: 

Wt.  of  magnesium  and  crucible =  g- 

Wt.  of  crucible = 

Wt.  of  magnesium 

Wt.  after  heating  (1) 

Wt.  after  heating  (2) 

(31) 


32  Laboratory  Experiments. 

Wt.  after  heating  (3) 

Gain  in  weight 

Per  cent,  of  gain 

Eq.  Wt.  of  magnesium 

Compare  your  result  with  the  values  obtained  by  two  other  students 
for  the  equivalent  weight  of  magnesium. 

Assuming  the  atomic  weights  of  magnesium  and  oxygen  (K.  81;  S. 
inside  rear  cover)  calculate  from  your  data  the  formula  for  the  oxide 
of  magnesium. 

16.  Equivalent  Weight  of  Mercury.     (Quant.)   (Optional).     (Two 
students  working  together.)     Dry  some  mercuric  oxide  by  heating  it 
lightly  above  100°.     Weigh  out  from  3  to  4  g.  of  the  oxide  in  a  hard 
glass  test  tube  at  least  15  cms.  long.     Support  the  tube  in  a  nearly 
horizontal  position  and  connect  it  air  tight  with  the  aspirator  used  in 
the  determination  of  the  weight  of  a  liter  of  oxygen,     Fig.  7.     Proceed 
as  in  the  decomposition  of  the  chlorate,  being  careful  however,  to  heat 
the  tube  as  little  as  possible  beyond  the  oxide.     Heat  till  the  oxide 
has  disappeared.     Determine  the  volume  of  the  oxygen,   noting  its 
temperature,  and  the  barometric  pressure.     Reduce  the  volume  of  the 
gas  to  standard  conditions  and  calculate  the  weight  of  the  oxygen. 
Calculate   also   the    percentage   composition   of    mercuric   oxide,    the 
equivalent  weight  of  mercury,  and  the  formula  of  the  oxide.     Write 
a  molecular  equation  for  the  chemical  change. 

17.  Composition  of  an  Oxide  of  Iron.     (Quant.) 

Weigh  an  evaporating  dish,  and  place  in  it  about  1.5-2.0  grams  of 
bright  iron  wire,  and  weigh  again.  Cover  the  dish  with  a  watch  glass, 
convex  side  down  (why  convex  side  down?)  and  add  20  cc.  of  dilute 
nitric  acid.  Heat  on  the  water  bath,  and  when  all  the  iron  has  dis- 
solved rinse  the  watch  glass  and  remove  it.  Evaporate  the  solution 
to  dryness  on  a  beaker  of  boiling  water.  Then  place  the  dish  on  a 
triangle  supported  on  a  ring-stand  and  heat  carefully  with  a  burner 
held  in  the  hand.  (Instructions.)  After  red  fumes  cease  to  be  given 
off,  heat  the  residue  to  redness.  Allow  to  cool  and  weigh.  Finally,  heat 
to  constant  weight. 

When  the  red  gases  are  set  free,  the  nitrate  of  iron  is  decomposing. 
The  red  fumes  show  the  presence  of  one  of  the  oxides  of  nitrogen  in  the 
gaseous  products.  The  residue  is  an  oxide  of  iron.  The  increase  in 
weight  represents  the  oxygen  combined  with  the  iron. 


Equivalent  Weights,  Formulas,  Equations.      33 


From  the  data  obtained,  calculate  the  equivalent  weight  of  iron. 

Assuming  the  atomic  weights  of  iron  and  oxygen,  calculate  from  your 
data  the  formula  for  the  oxide  of  iron.  What  is  the  name  of  this  oxide? 

Magnesium,  zinc,  tin  or  copper,  may  be  used  instead  of  iron,  and  the 
method  of  operation  is  essentially  the  same.  The  residues,  however, 
in  some  cases,  cannot  be  dried  on  the  water  bath.  Extra  caution  must 
therefore  be  exercised  in  heating  them  to  avoid  loss  by  spattering. 

N.  B.  Whilst  the  solution  is  evaporating  the  next  experiment  may 
be  performed.  Regulate  the  heating  by  lowering  the  flame  of  the  burn- 
er so  that  the  water  boils  off  slowly,  and  then  proceed  with  other 
work.  Waiting  around  for  evaporations  may  be  good  training  in  the 
cultivation  of  patience,  but  evidently  such  time  is  not  to  be  credited 
as  laboratory  work. 

18.  Equivalent  Weight  of  Zinc  by  Displacing  Hydrogen.  (Quant.) 
(Two  students  working  together.) 

a.  First  fill  the  pneumatic  trough  and  1-1.  bottle  with  water  so 
that  the  liquid  may  acquire  the  temperature  of  the  room.  Fit  a  100 
cc.  flask  with  a  drop-funnel  and  a  delivery  tube  as  shown  in  Fig.  9. 
Test  the  apparatus  to  see  that  it  is  air-tight.  Weigh  a  piece  of  pure 


Fig.  9. 

zinc  (Storeroom)  of  about  2  g.     Without  detaching  the  platinum  wire 
from  the  glass  rod,  wrap  it  around  the  zinc  (why?)  and  slip  the  whole 
Chem.— 3 


34  Laboratory  Experiments. 

into  the  flask.  Now  fill  the  apparatus  completely  from  the  stop-cock 
of  the  drop-funnel  to  the  tip  of  the  delivery  tube  with  water.  Invert 
the  liter  bottle  with  water  in  the  pneumatic  trough,  and  slip  the  end  of 
the  delivery  tube  into  the  mouth  of  the  bottle. 

Fill  the  bulb  of  the  funnel  with  pure  concentrated  hydrochloric  acid 
and  admit  this  to  the  flask,  in  such  a  way  that  a  steady  stream  of  gas 
flows  into  the  collecting  bottle.  When  the  metal  is  entirely  dissolved, 
drive  all  gas  over  into  the  bottle  by  pouring  water  through  the  funnel. 

The  weight  of  the  hydrogen  displaced  in  this  experiment  is  not  to 
be  ascertained  by  direct  weighing,  but  by  determining  the  volume  of 
the  gas  and  calculating  its  weight  from  this  and  the  density  of  the  gas. 
Proceed  as  follows:  After  the  gas  has  assumed  room  temperature, 
equalize  the  levels  of  the  water  inside  and  outside  the  bottle.  While 
in  this  position  cork  the  bottle  and  remove  it  from  the  trough.  Find 
the  volume  of  the  gas.  (See  Exp.  2c.)  Record  the  temperature  of 
the  water  in  the  bottle.  Read  the  barometer  and  the  temperature 
near  it.  Since  the  hydrogen  is  mixed  with  water  vapor,  correct  for 
the  latter  (Exp.  10)  and  then  reduce  the  volume  of  the  hydrogen,  by 
rule,  to  standard  conditions.  From  the  data  obtained,  calculate  the 
equivalent  weight  of  zinc,  i.  e.,  the  quantity  of  zinc  which  displaces 
1.008  g.  of  hydrogen: 

Wt.  of  hydrogen     :     Wt.  of  zinc     =     1.008     :     x. 

What  is  the  valence  of  zinc? 

How  many  molecular  weights  of  hydrogen  chloride  are  required  to 
furnish  the  hydrogen  that  one  atomic  weight  of  zinc  will  displace? 
Write  the  part  of  the  equation  that  expresses  your  answer  to  this 
question. 

b.  Determine  the  hydrogen  equivalent  of  some  other  metal  by  the 
above  method, — magensium,  using  0.7  to  1.0  g.;  aluminum,  0.5  to  0.8 
g.;  or  iron,  1.5  to  2.0  g.  Whilst  the  zinc  is  dissolving,  clean  and  weigh 
out  the  other  metals. 

Refer  to  the  experiment,  "Weight  of  a  Liter  of  Oxygen"  and  use  the 
plan  outlined  there  in  recording  your  experimental  data. 

19.     Combining  Weights  of  Zinc  and  Chlorine.     (Quant.) 

Weight  an  evaporating  dish  and  place  in  it  about  2  g.  of  pure  zinc 

(Storeroom)  and  weigh  again.       Wrap  the  platinum  wire  around  the 

zinc,  and  then  dissolve  it  in  diluted  hydrochloric  acid  (1:2),  keeping 

the  dish  covered  with  a  watch  glass  during  the  operation.     Remove 


Equivalent  Weights,  Formulas,  Equations.      35 

the  glass  and  wire,  rinse  both,  and  evaporate  the  solution  on  the  water 
bath  as  far  as  possible.  Place  the  dish  next  on  a  triangle  and  evaporate 
it  slowly  to  dryness.  Then  heat  the  white  residue  to  the  point  where 
it  has  melted  and  no  further.  When  the  dish  has  cooled  so  that  it 
can  be  borne  on  the  hand,  weigh  it  quickly.  Repeat  the  melting, 
cooling  and  weighing,  and  take  the  lower  result  as  correct. 

Calculate,  from  the  data  obtained,  how  much  chlorine  combines 
with  the  equivalent  weight  of  zinc  found  in  the  previous  experiment. 
This  amount  is  the  equivalent  weight  of  chlorine,  x, — 

Wt.  of  zinc     :     Wt.  of  chlor.     =     Equiv.  of  zinc     :     x. 

Assuming  the  atomic  weights  of  zinc  and  chlorine,  determine  the 
formula  of  the  compound. 

Express  the  whole  reaction  of  zinc  and  hydrochloric  acid  by  making 
the  equation  in  accordance  with  the  results  obtained  in  this  and  Exp. 
18.  Which  of  the  factors  in  the  equation  have  you  determined  ex- 
perimentally, and  which  not?  What  law  do  we  use  in  assuming  that 
the  undetermined  factors  are  correct? 

Record  your  data  in  tabular  form. 

20.  Law  of  Multiple  Proportions.  (Quant.)  (Two  students 
working  together.) 

Dry  samples  of  potassium  chlorate  and  potassium  perchlorate  on 
watch  glasses  by  gently  warming  on  a  sand  bath,  or  radiator.  Weigh 
two  hard  glass  test  tubes,  and  introduce  about  1  g.  of  the  chlorate  in 
one  tube  and  weigh  again.  In  the  other  tube  place  about  the  same 
quantity  of  perchlorate  and  weigh.  Now  heat  the  tubes  cautiously 
and  decompose  the  compounds,  allowing  the  oxygen  to  escape.  Should 
white  vapors — a  cloud  of  solid  particles — appear  at  the  mouth  of  the 
tube  the  heating  must  be  slackened  for  a  time.  When  the  decomposi- 
tion is  complete  (test?)  allow  the  tubes  to  cool  and  weigh  the  tubes 
with  their  residues. 

Record  your  weighings  as  suggested  by  the  following  outline: 

Chlorate.     Perchlorate. 

Weight  of  tube  +  salt g.  g. 

Weight  of  tube 

Weight  of  salt  taken 

Weight  of  tube  +  residue g.  g. 

Weight  of  tube '.  . .  . . 

Weight  of  residue 

Loss  in  weight  of  salt 


36  Laboratory  Experiments. 


Assuming  that  the  loss  in  weight  in  these  experiments  represents 
the  oxygen  in  the  compound,  calculate  how  much  oxygen  is  combined 
with  one  gram  of  each  of  the  residues.  Assuming  that  the  residues 
are  the  same  substance,  potassium  chloride,  show  how  the  data  just 
computed  illustrate  the  law  of  multiple  proportions. 

Assuming  that  the  formula  for  potassium  chloride  is  KC1,  and  its 
molecular  weight  74.6,  calculate  the  quantities  of  oxygen  combined 
with  74.6  parts  of  the  chloride  in  each  of  the  above  experiments: 

Wt.  of  chloride     :     Wt.  of  oxygen     =     74.6     :     x. 

Where  x  is  the  weight  of  oxygen  combined  with  one  molecular  weight 
of  potassium  chloride  in  each  experiment. 

Taking  the  atomic  weight  of  oxygen  16,  deduce  the  formula  for  each 
of  the  substances  from  the  data  now  in  hand,  and  finally  write  equa- 
tions that  will  express  the  quantitative  relations  here  determined. 

21.  Law  of  Multiple  Proportions.  (Quant.,  optional,  two  students 
working  together.) 

Fit  a  piece  of  hard  glass  tubing  at  least  20  cm.  long  with  glass  tubes 
and  corks  as  shown  in  Fig.  6,  and  mount  the  tube  for  heating  with  the 
Bunsen  flame.  Make  sure  that  the  apparatus  is  air-tight.  On  the 
radiator  dry  specimens  of  pure  litharge  and  lead  dioxide.  Weight 
two  porcelain  boats  and  fill  the  boats  with  litharge  and  the  dioxide 
respectively,  using  at  least  2  g.  of  each,  and  then  weigh  again. 

Place  the  boats  in  the  tube  of  hard  glass  so  that  the  points  of  the 
boats  touch  in  the  center  of  the  tube.  Connect  the  tube  with  a  supply 
of  dry  hydrogen  or  illuminating  gas  and  pass  a  gentle  stream  of  the  gas 
through  the  tube.  Test  the  issuing  gas  to  see  it  is  free  from  air. 
Now  heat  the  boats  moderately  waving  the  flame  to  and  fro  at  first. 
What  collects  in  the  cooler  part  of  the  tube?  Where  does  it  come 
from?  When  minute  metallic  globules  appear  at  any  point  in  the  boats , 
withdraw  the  heat  from  that  part  of  the  tube  for  a  time.  The  residue 
is  more  easily  removed  from  the  boats  afterward  if  it  does  not  fuse. 
After  ten  minutes  of  heating,  remove  the  flame  and  allow  the  boats 
to  cool  in  a  stream  of  hydrogen.  Remove  the  boats  and  weigh  them. 
Reheat  then  until  constant  weight  is  obtained. 

Record  the  data  so  that  the  weighings  for  the  oxides  will  appear  in 
parallel  columns. 

The  loss  in  weight  represents  the  oxygen  removed  from  the  oxides. 
Calculate  the  quantities  of  lead  in  combination  with  8  g.  of  oxygen  in 


Equivalent  Weights,  Formulas,  Equations.      37 


the  oxides  and  apply  the  law  of  multiple  proportions  to  the  data  thus 
obtained. 

22.     Law  of  Dulong  and  Petit.     Formulate  the  law. 

Take  the  equivalent  weights  you  have  found  experimentally  in 
this  chapter,  viz.:  magnesium,  mercury,  zinc,  iron,  aluminum,  etc.,  and 
multiply  each  by  the  corresponding  specific  heat.  (K.  78;  S.  135.) 
If  the  product  is  about  6.4,  the  atomic  weight  is  the  same  as  the 
equivalent  weight.  If  not,  then  the  smallest  multiple  that  will  bring 
the  product  up  to  about  6.4  is  the  valence  of  the  element,  and  the 
product  of  this  integer  and  the  equivalent  weight  is  the  atomic  weight 
of  the  element. 

Arrange  your  data  in  the  form  of  a  table  under  the  following  heads: 


i 

X 

O 

^   +j 

u 

.a     -M 

.a 

3 

.£  -o 

3  C 

<U     rt 
C,    V 
C/3  J3 

«! 

"c3 

If 

<  > 

—•     c 
<  JS 

Plot  the  specific  heats  (K.  78,  or  S.  135),  against  the  atomic  weights 
and  name  the  curve  obtained  by  connecting  the  points. 

Compare  this  curve  with  the  one  obtained  by  plotting  the  pres- 
sures against  corresponding  volumes  for  a  given  volume  of  gas,  the 
temperature  remaining  constant  (Boyle's  Law).  What  is  the  equa- 
tion for  the  curve?  A  suggestion: — Assume  that  you  have  100  cc. 
of  air  at  constant  temperature  and  under  a  pressure  of  one  atmos- 
phere. The  volume  corresponding  to  pressures  of  2,  4,  5,  8,  and  10 
atmospheres  are  easily  deduced  by  applying  Boyle'  slaw.  The  vol- 
umes corresponding  to  pressures  of  0.5,  0.2,  and  0.1  atmosphere 
should  be  included,  and  the  different  pressures  and  corresponding 
volumes  tabulated,  as  well  as  the  products,  P  X  V. 

What  is  the  weight  of  a  column  of  mercury  76  cms.  high  and  1 
sq.  cm.  in  cross  section?  The  density  of  mercury  is  13.59.  What 
is  the  pressure  of  one  atmosphere  in  grams? 


232445 


CHAPTER  VI. 


REVIEW  EXERCISES. 

a.  Referring  to  your  laboratory  notebook,  consider  the  chemical 
changes  involved  in  the  experiments  that  you  have  performed,  and 
write   equations   for  the   reactions.     In   several    of   the   experiments 
nitric  acid  was  employed,  and  since  the  reactions  of  metals  and  nitric 
acid  are  complex,  the  nature  of  the  products  depending  upon  the  con- 
centration of  the  acid  and  the  temperature,  the  equations  for  these 
changes  need  not  be  attempted. 

b.  What   exceptions   do   your  experiments   furnish   to  the   state- 
ment that  hydrogen  is  obtained  by  the  reaction  of  acids  with  certain 
metals,  e.  g.,  zinc  or  magnesium. 

c.  Experiment  11  is  often  cited  as  an  illustration  of  catalytic  ac- 
tion, and  manganese  dioxide  is  referred  to  as  a  catalytic  agent.     In 
what  other  experiments  did  you  employ  catalytic  agents? 

d.  Calculate  the    molar    solubility  of  potassium  dichromate  from 
the  data  of  experiment  14. 

e.  What  is  the  mineralogical  name  for  the  oxide  of  iron  obtained 
in  experiment  17? 

f.  Give  a  reason  for  weighing  the  sample  of  zinc  chloride  obtained 
in  experiment  19  while  it  was  still  warm. 

g.  Compare  the  volume  of  the  oxygen  as  measured  when  stand- 
ing over  water  as  in  experiment   10,  Weight  of  a  Liter  of  Oxygen, 
with  its  volume  if  measured  over  mercury  under  the  same  conditions 
of  temperature  and  pressure. 

h.  How  much  zinc  chloride  will  be  furnished  by  dissolving  25  g. 
of  pure  zinc  in  pure  hydrochloric  acid?  Use  the  data  that  you  ob- 
tained in  experiment  18.  Calculate  the  quantity  of  the  chloride 
from  25  g.  of  zinc  using  the  equation  that  you  have  written. 

i.  What  chemical  test  would  establish  the  fact  that  the  residue 
left  when  potassium  chlorate  and  perchlorate  are  decomposed,  ex- 
periment 20,  consists  of  a  chloride? 

(38) 


Review  Exercises.  38 

j.  Criticize  the  statement  that  the  amount  of  oxygen  in  lead  di- 
oxide is  twice  that  of  the  monoxide.  (Experiment  21.) 

k.  Calculate  the  number  of  calories  required  to  raise  the  temper- 
ature of  207  g.  of  lead  1  .  What  term  in  the  table  of  exercise 
22  properly  designates  this  value? 

1.  Assuming  the  molecular  weights  of  cupric  sulphate  and  water, 
deduce  a  formula  for  blue  vitriol  from  the  date  of  experiment  13d. 

m.     Make  a  similar  calculation  with  the  data  of  experiment  13c. 


CHAPTER  VII. 


THE  HALOGENS. 

Note  8.  Read  carefully  the  direction  before  beginning  an  experi- 
ment. The  absence  of  italics  and  black-faced  type  means  that  every 
line  is  significant.  If  you  grasp  the  purpose  of  the  experiment  at  the 
outset  and  then  follow  the  outline  thoughtfully  your  daily  achieve- 
ments will  make  steady  gains.  Hasty  work  and  makeshift  mount- 
ing of  apparatus  generally  necessitate  a  repetition  of  the  work,  be- 
sides causing  much  trouble  and  needless  nervous  strain. 

The  bringing  together  of  substances  as  directed  so  that  chemical 
changes  take  place  is  in  itself  of  little  cultural  value.  Most  of  the 
manual  occupations  are  quite  as  complex  and  require  fully  as  much 
thought  and  planning.  A  notebook  record  giving  the  results  of  ex- 
periments in  such  terms  as — dense  red  fumes  arose,  a  bright  flame 
was  seen,  a  white  precipitate  formed,  a  pop  and  an  explosion  followed, 
a  pungent  odor  was  given  off,  and  linked  with  a  "hence"  or  a  "there- 
fore" evidently  shows  that  the  inwardness  of  the  phenomena  still 
remains  practically  untouched. 

"A  primrose  by  the  river's  brim 
A  yellow  primrose  was  to  him 
And  nothing  more." 

The  exercises  call  for  reflection  and  require  prolonged  concentration 
of  attention.  What  is  the  nature  of  the  changes  presented  in  the  ex- 
periment? Write  equations  for  the  actions.  What  underlying  prin- 
ciple is  illustrated?  Give  your  interpretation  of  the  phenomena.  Is 
the  experiment  related  to  some  previous  one  or  to  a  class-room  illus- 
tration? Point  out  how  your  results  differ  from  the  statements  of 
the  text.  How  do  you  account  for  such  discrepancies?  Do  not 
stop  with  knowing,  but  keep  on  doing  until  the  knowing  becomes  a 
habit  ever  yours. 

23.  Preparation  of  the  Hydrogen  Halides.  (Two  students  working 
together.) 

a.  Place  small  quantities  of  each  of  the  following  substances  in 
separate  test  tubes:  Fluorspar,  common  salt,  sodium  or  potassium  bro- 

(40) 


The  Halogens.  41 


mide,  sodium  or  potassium  iodide.  Add  several  drops  of  concentrated 
sulphuric  acid  to  each.  Describe  what  happens  in  each  case.  Blow- 
moist  air  across  the  mouth  of  the  tubes.  (?)  Hold  a  piece  of  blue 
litmus  paper  near  the  mouth  of  each  tube.  (?)  Hold  a  strip  of  filter 
paper  dipped  in  starch  paste  in  the  mouth  of  each  tube.  (?)  Bring 
a  rod  dipped  in  ammonium  hydroxide  near  the  mouth  of  each  test  tube 
and  gradually  lower  it  into  the  tube.  (?)  Try  the  effect  of  heating 
if  the  action  in  the  cold  is  slight.  (?)  What  evidence  is  there  that  the 
gases  evolved  in  some  of  the  interactions  consist  of  two  or  more  gas- 
eous produce?  In  the  reactions  where  you  suspect  a  mixture  of  gas- 
eous products  try  to  identify  some  of  those  not  yet  recognized. 
Remove  the  residues  from  the  tubes,  and  examine  the  tubes.  Do 
any  of  them  show  signs  of  corrosion.  (?) 

b.  Perform  another  series  of  experiments  using  the  same  salts,  but 
substitute  concentrated  phosphoric  acid  for  sulphuric.     Compare  the  re 
suits  with  those  noted  in  the  trials  with  sulphuric  acid. 

c.  Arrange  the  halogen  hydrides  in  the  order  of  their  stability, 
and  compare  this  order  with  their  heats  of  formation.     (K.  284.) 

What  facts  gathered  here  support  the  statement  that  hydrogen 
iodide  is  a  reducing  agent?  (K.  116;  S.  167.) 

Beginning  with  this  chapter  include  in  your  record  equations  for 
the  interactions  studied. 

Which  of  the  foregoing  methods  of  preparation  is  evidently  not 
suited  in  obtaining  the  halogen  hydride  in  a  fairly  pure  state? 

d.  In  separate  test  tubes  place  abou't  1  g.  of  each  of  the  following 
chlorides:     Sodium  chloride,  ammonium  chloride,  mercuric  chloride, 
and  calcium  chloride.       Add    a    few  cc.    of    concentrated    sulphuric 
acid     to    each    chloride,    and    test    for    the    presence    of    hydrogen 
chloride  at   the    mouth   of  the   test   tube  by  one  of  the  methods  em 
ployed  in  a.     Heat  the  contents  of  the  tubes  if  no  action  occurs  at 
room  temperatures.     (?) 

24.  Preparation  of  the  Halogens.  (Two  students  working  to- 
gether.) 

a.  In  separate  test  tube  place  about  a  gram  of  each  of  the  following 
substances: — Manganese  dioxide,  cupric  oxide,  lead  dioxide,  litharge,, 
potassium  chlorate,  potassium  permanganate.  Prepare  some  test 
papers,  by  dipping  strips  of  filter  paper  in  starch  paste  to  which  you 
have  added  2-3  drops  of  potassium  iodide  solution. 


42  Laboratory  Experiments. 


Add  a  few  drops  of  concentrated  hydrochloric  acid  to  the  first  test 
tube,  and  note  the  color  (?)  and  odor  (?)  of  the  gas.  If  no  action  takes 
place  in  the  cold,  heat  gently.  Dip  into  the  gas  one  of  the  test  papers. 
(?)  Suspend  one  of  the  test  papers  in  the  gas  for  a  minute  or  two.  (?) 
Empty  the  contents  of  the  tube  or  set  it  aside;  then  treat  the  contents 
of  the  second  tube  with  hydrochloric  acid  and  test  the  gas  as  before. 
(?)  Experiment  with  the  other  substances  in  the  same  way,  and  note 
carefully  differences  in  behavior.  How  do  you  account  for  the  differ- 
ences? Do  not  heat  the  tube  containing  the  potassium  permanganate. 

b.  Powder  about  1  g.  of  sodium  chloride  and  mix  it  with  about 
2  g.  of  pulverized  manganese  dioxide.  Place  the  mixture  in  a  test 
tube,  and  add  2-3  cc.  of  sulphuric  acid  (1  Acid  :  1  Aq.)  In  mixing 
sulphuric  acid  and  water  add  the  acid  cautiously  to  the  water.  Stir 
up  the  contents  of  the  test  tube,  and  heat  gently,  and  identify  the  gas. 
(?)  What  other  substances  might  be  substituted  for  the  common  salt 
in  this  experiment?  (Refer  to  Exp.  23d.)  For  the  manganese  dioxide. 
(?) 

Modify  the  above  experiment  by  using  potassium  bromide  in  place 
of  soduim  chloride.  (?) 

Finally  try  the  experiment  using  posassium  iodide  instead  of  potas- 
sium bromide.  (?) 

25.  Properties  of  Iodine  and  Bromine.  Place  20  cc.  of  water  in 
each  of  two  test  tubes.  Add  a  drop  of  bromine  to  one  test  tube  and 
a  few  crystals  of  iodine  to  the  other.  Shake  the  tubes  to  hasten. 
solution.  Care: — Do  not  spill  bromine  on  the  hands.  Iodine  stains 
may  be  removed  by  washing  with  hypo,  a  solution  of  sodium  thio- 
sulphate. 

Divide  the  solution  of  each  tube  between  four  test  tubes.  Add  2 
cc.  of  ether  to  one  of  the  samples  of  bromine  water,  and  to  one  of 
the  iodine  solutions.  After  shaking  the  tubes  note  the  relative 
solubility  in  water  and  ether  as  shown  by  the  depth  of  color  of  the 
layers. 

To  another  pair  of  solutions  add  2  cc.  of  carbon  disulphide  and 
shake.  (?) 

To  the  third  pair  add  chloroform.     (?) 
To  the  fourth  pair  add  10-15  cc.  of  starch  paste.     (?) 
To  a  few  crystals  of  iodine  add  5  cc.  of  water  and  shake.     Now  add 
a  small  crystal  of  potassium  iodide  and  shake  again.     (?)     Interpret. 


The  Halogens.  43 


What  is  tincture  of  iodine? 

26.  Preparation  of  Hydrobromic    Acid.      (Optional;  two  students 
working  together.) 

Set  up  the  apparatus  figured  on  p.  108  K.  or  S.  p.  162.  In  the  flask, 
250  cc.  capacity,  place  about  5  g.  of  red  phosphorus  mixed  with  an 
equal  volume  of  sand  (use  of  this?)  and  5  cc.  of  water.  In  the  U-tube 
place  broken  glass  mixed  with  a  little  red  phosphorus  (use  of  this?). 
Pour  into  the  drop  funnel  10  cc.  of  bromine,  using  a  funnel  in  trans- 
ferring the  bromine  to  the  drop  funnel.  Extreme  care  should  be  used 
in  handling  bromine.  To  start  the  reaction,  allow  the  bromine  to 
trickle  slowly  on  the  phosphorus.  Collect  the  gas  by  dissolving  it  in 
water  so  arranged  that  the  solution  can  not  suck  into  the  connecting 
U-tube  and  generator. 

Try  the  effect  of  moist  air  on  the  issuing  gas.  (?)  Hold  in  the  gas 
a  glass  rod  dipped  in  ammonium  hydroxide.  (?)  Test  its  action  on 
moist  litmus  paper.  (?) 

Test  the  action  of  a  portion  of  the  acid  on  zinc.  (?)  Mix  another 
portion  with  manganese  dioxide  and  heat  gently.  (?)  Add  a  few 
drops  of  chlorine  water  to  another  portion.  (?)  Taste  a  drop  of  the 
diluted  solution.  To  another  portion  add  a  few  drops  of  silver  nitrate 
solution,  and  expose  the  precipitate  to  sunlight.  (?) 

If  water  were  allowed  to  trickle  from  the  drop  funnel  of  the  apparatus 
what  substance  would  be  suitable  to  have  in  the  flask  instead  of  the 
phosphorus  and  sand  mixture  for  the  preparation  of  hydrogen  bromide? 

What  advantage  has  this  method  of  preparing  hydrobromic  acid 
over  the  interaction  of  sulphuric  acid  and  a  bromide?  (Cf.  Exp.  23a.) 

27.  Preparation     of     Hydriodic     Acid.     (Hood;     Optional;     Two 
students  working  together.) 

Place  10  g.  of  powdered  iodine  in  a  flask  containing  about  75  cc. 
of  water,  and  provided  with  a  one-hole  cork  and  short  glass  tube. 
Pass  hydrogen  sulphide  into  the  mixture  of  iodine  and  water.  Use 
the  same  form  of  generator  as  was  used  in  preparing  hydrogen.  For 
preparing  hydrogen  sulphide  use  ferrous  sulphide  in  place  of  zinc.  At 
first  the  cork  on  the  flask  containing  the  iodine  and  water  must  be 
loosened  occasionally  to  allow  the  air  to  be  displaced  by  the  hydrogen 
sulphide.  Shake  the  iodine  and  water  mixture  constantly  to  hasten 
the  process.  Describe  what  happens.  Continue  passing  the  gas  until 


44 


Laboratory  Experiments. 


the  solution  no  longer  becomes  brown  on  shaking.     Warm  and  filter 
the  solution.     What  is  the  residue  on  the  filter? 

Distil  the  filtrate  fractionally,  figure  10,  collecting  first  the  portion 
that  comes  over  between  99  and  100,  and  then  the  fractions  boiling 


Fig.  10. 

between  100-105,  105-110,  and  so  on.  Use  a  small  flame  and  place 
a  gauze  under  the  distilling  flask.  Stop  when  the  liquid  is  nearly  all 
distilled  off.  Note  the  highest  temperature  reached.  Pour  the  liquid 
left  in  the  flask  into  a  test  tube  and  keep  the  series  of  distillates  for 
the  following  experiments: 

Add  silver  nitrate  solution  to  the  different  fractions,  using  only  a 
part  of  the  solution  in  the  case  of  the  two  highest  boiling  fractions. 
At  what  temperature  did  the  most  concentrated  acid  distil  over? 

Try  the  action  of  a  portion  of  the  strong  acid  on  zinc.  (?)  Test 
with  litmus  paper  the  higher  boiling  fractions.  (?)  Warm  some  of 
the  concentrated  acid  with  a  little  powdered  manganese  dioxide? 

What  substance  causes  the  color  of  the  higher  fractions  and  of  the 
residue?  Test  your  conclusion  by  an  experiment. 

Record  the  maximum  boiling  temperatures  of  aqueous  solutions  of 
hydrogen  chloride,  hydrogen  bromide  and  hydrogen  iodide  and  the 
corresponding  concentrations  of  these  mixtures.  (K.  50,  108,  115:  S. 
120,  163,  167.) 


The  Halogens. 


45 


28.     Properties    of    Chlorine.     (Hood.)     (Two    students    working 
together.) 

a.     Fit   up  a  250  cc.  flask  with  a  drop  funnel  and  delivery  tube 

as  shown  in  Fig.  11.  Since 
chlorine  gas  destroys  rubber 
tubing  make  the  rubber  con- 
nections short.  Test  the  ap- 
paratus to  see  that  it  is  air- 
tight. Place  in  the  flask  about 
20  g.  of  potassium  perman- 
ganate, and  from  the  drop 
funnel  admit  drop  by  drop 
hydrochloric  acid  (1:1).  Reg- 
ulate the  flow  so  that  too  rapid 
generation  of  gas  is  not  pro- 
duced. When  the  air  has  all 
been  displaced  from  the  gene- 
rator (how  judge?)  fill  four 
Cover  the  bottles  with 


Fig.  11. 

dry,    wide-mouthed    bottles   with   the   gas. 
glass  plates  as  soon  as  they  are  full. 

b.  When  all  the  bottles  have  been  filled  pass  a  few  bubbles  of  gas 
into  5  cc.  of  solutions  of  potassium  bromide  and  potassium  iodide  in 
separate  test  tubes.     (?)     Add  a  few  drops  of  carbon  disulphide  to 
each    test    tube    and    shake.     (?)     Interpret    the    results.     Pass    the 
superfluous  gas  into  a  solution  of  sodium  hydroxide. 

Dismantle  the  chlorine  generator  when  sufficient  gas  has  been  col- 
lected, pour  the  contents  of  the  generator  into  the  sink  and  flush  with 
water. 

c.  Fill  a  wide-mouth  bottle  with  hydrogen  and  place  it  mouth  to 
mouth  with  a  bottle  of  chlorine.     Mix  the  gases  by  inverting  a  few 
times  but  do  not  conduct  this  operation  in  direct  sunlight.     Hold  the 
mouth  of  each  bottle  near  a   Bunsen  flame.     (?)      Blow  the  breath 
into  one  of  the  bottles  (?),  and  to  the  other  add  a  few  drops  of  water; 
shake,  and  test  the  solution  with  litmus  paper.     (?) 

d.  In  one  of  the  bottles  of  gas  place  moist  red  and  blue  litmus  paper; 
a  paper  on  which  is  ordinary  writing  and  also  printing;  a  few  blades  of 
grass,  and  a  piece  of  colored  calico.      (?) 


46  Laboratory  Experiments. 

e  In  another  bottle  of  gas  scatter  a  pinch  of  finely  powdered  an- 
timony. (?)  (K.  325;  S.  468.) 

f.  Connect  a  glass  nozzle  with  the  illuminating  gas  supply,  and  lower 
a  small,  burning  gas  flame  into  the  fourth  bottle.  (?)  Blow  the  breath 
gently  into  the  bottle  after  withdrawing  the  jet.  Interpret  the  results. 

29.  Identification  of  Halogen  Compounds.  If  you  were  given  four 
white  substances  and  told  that  they  were  the  fluoride,  chloride,  bromide 
and  iodide  of  some  metal,  state  what  experiments  you  would  make  in 
order  to  identify  the  halogen  constituent  of  each.  Negative  results, 
showing  that  one  is  not  a  chloride,  bromide,  or  iodide,  and  is  therefore 
a  fluoride,  must  be  confirmed  by  an  experimental  test. 

Ask  the  instructor  for  an  unknown  halogen  compound  and  then 
proceed  to  identify  it,  recording  the  different  tests  that  you  conducted  . 
Report  your  conclusion  to  the  instructor. 


CHAPTER   VIII. 


ACIDS,     BASES,     SALTS.     CHEMICAL     EQUILIBRIUM. 

30.     Properties  of  Acids  and  Bases. 

a.  Prepare  a  sample  of  dilute  hydrochloric  acid  by  adding  5  cc. 
of  the  concentrated  acid  to  25  cc.  of  distilled  water,  and  examine  the 
solution  in  respect  to  taste,  behavior  toward  litmus,  phenolphthalein 
(use  but  f  or  2  drops),  action  on  a  bit  of  magnesium  ribbon  or  zinc, 
action  on  baking  soda  or  marble.     Add  one  drop  of  concentrated  acid 
to  10  cc.    of   water  and  see  whether  the  solution  tastes  sour.     De- 
termine the  dilution  at  which  you  can  barely  detect  the  sour  taste  of 
the  acid.     (?)      Test  the  very  diluteso  lution  with  litmus.     (?) 

Test  a  sample  of  dilute  nitric  acid  (1:  10)  with  the  same  reagents 
used  in  determining  the  characteristics  of  hydrochloric  acid.  (?) 

The  above  properties  are  shown  by  other  acids. 

What  element  (or  elements)  is  found  in  the  formulas  of  all  the  acids 
on  the  laboratory  shelf  of  liquids?  If  HnA  is  adopted  as  the  general 
formula  for  acids,  where  n  may  be  1,  or  2,  or  3,  etc.,  underscore  the 
radical  represented  by  A  in  the  following  acids, — sulphuric,  carbonic, 
nitric,  chloric,  hydriodic,  perchloric,  sulphurous,  phosphoric.  Indicate 
the  valence  of  the  radicals  underscored. 

b.  Disslove  5  g.  of  lye  in  50  cc.  of  water,  and  test  the  solution  with 
the  same  reagents  employed  in  a.     Repeat  the  experiments  with  a 
sample  of  lime  water.     Write  the  chemical  name  and  formula  for, — 
lime  water,  lye,  caustic  potash.     Other  bases  have  the  properties  com- 
mon to  caustic  soda.     What  radical  is  common  to  bases?     Write  a 
general  formula  for  bases. 

Tabulate  the  observations  in  a  and  b  under  the  following  heads: — 
Taste,  litmus  test,  phenolphthalein  test,  action  on  magnesium,  action 
on  baking  soda. 

c.  In  separate  test  tubes,  or  watch  glasses,  place  about  1  cc.  of  the 
following  solutions  and  note  their  behavior  toward  litmus: — borax, 
ferric  chloride,  magnesium  sulphate,  sodium  acid  carbonate,  sodium 
carbonate,  sodium  acid  sulphate,  sodium  chloride.     (?)     How  do  you 

(47) 


48  Laboratory  Experiments. 

account  for  the  differences  in  behavior  of  the  solutions  toward  litmus? 
(K.  130;  S.  163,  353.) 

29.     Neutralization.     (Two  students  working  together.) 

a.  Mount  a  pair  of  clean  burettes  fitted  with  nozzles.     (Exp.  4.) 
Prepare  about  400  cc.   of  sodium   hydroxide  solution  by  dissolving 
20  g.  of  the  solid.     Prepare  also  about  200  cc.  of  dilute  hydrochloric 
acid  by  adding  20  cc.  of  concentrated  acid  to  200  cc.  of  water.     Shake 
the  solutions  before  using.      Fill  one  of  the  burrettes  with  base  and  the 
other  with  acid.     Now,  into  a  small  beaker  or  flask  run  10  cc.  of  the 
acid  and  add  two  drops  of  phenolphthalein.     Place  the  vessel  under 
the  other  burette  and  allow  the  alkali  to  run  into  the  acid  slowly, 
stirring  constantly,  until  a  drop  produces  a  perceptible  pink  tint  in 
the  whole  solution.     If  too  much  alkali  is  accidentally  added,  run  in  a 
drop  or  two  of  acid  to  discharge  the  pink  color,  and  then  add  alkali 
again    until    the    solution    becomes    faintly    pink.     Finally    read    the 
volumes  of  acid  and  base  used.     Make  a  second  and  third  trial  with 
fresh  quantities  of  acid  using  15-20  cc.  in  each  determination.     Read 
carefully  the  volumes  of  acid  and  base  used,  and  calculate  the  ratio 
of  the  volumes  in  each  titration.     Arrange  your  data  in  the  form  of  a 
table  under  the  following  heads: — Determination  number,  volume  of 
alkali,  cc.,  volume  of  acid,  cc.,  indicator,  ratio  of  base  to  acid.     In- 
terpret the  ratios. 

Find  the  average  of  your  ratios  and  express  the  percentage  difference 
between  the  mean  value  and  the  single  determination  differing  most 
from  it.  Compute  the  volume  of  alkali  that  would  be  required  to 
neutralize  25  cc.  of  your  acid. 

b.  (Quant.)     Neutralize  exactly  25  cc.  of  the  acid  with  your  solution 
of  alkali,  and  note  the  volume  of  alkali  required.     Transfer  the  solu- 
tion to  a  weighed  evaporating  dish,   rinse  the  vessel,  and  add  the 
rinsings  to  the  main  solution.     Evaporate  to  dryness  on  the  water 
bath,  and  finally  heat  the  residue  gently  over  a  small  Bunsen  flame. 
Allow  it  to  cool  and  then  weigh.     What  is  the  residue? 

Assuming  the  equation  for  the  neutralization  (K.  122;  S.  241),  cal- 
culate from  the  weight  of  the  residue  the  number  of  grams  of  hydro- 
gen chloride  and  sodium  hydroxide  per  liter  of  your  solutions.  Ex- 
press the  concentration  (strength)  of  your  solutions  in  terms  of  nor- 
mal. 


Acids,  Bases,  Salts.  49 

c.  (Quant.)     Obtain  a  sample  of  standard  sulphuric  acid  from  the 
instructor.     Determine  the  concentration  of  your  alkali  solution  by 
titration  against  the  standard   sulphuric  acid.     Make   at   least  two 
determinations,  and  do  not  use  less  than  10  cc.  of  acid  for  each.     Ex- 
press the  concentration  of  your  alkali  in  terms  of  normal.     Compare 
this  value  with  that  obtained  in  b,  and  express  the  percentage  difffer- 
ence.     Submit  the  result  to  the  instructor  for  criticism. 

d.  (Quant.)     Secure  from  the  instructor  a  sample  of  an  acid,  an 
unknown,  and  determine  its  concentration  by  titration  against  your 
standard  alkali,  and  report  the  value  to  the  instructor. 

What  weight  of  ammonium  chloride  will  result  from  the  neutral- 
ization of  50  cc.  of  fifth  normal  ammonium  hydroxide?  What  vol- 
ume of  fourth  normal  (  N/4  )  acid  will  be  required? 

32.  Double  Decomposition.  Precipitation.  (Two  students  working 
together.) 

Place  the  following  in  separate  test  tubes, — dilute  hydrochloric 
acid,  potassium  chlorate  solution,  potassium  chloride  solution,  a 
solution  of  some  other  chloride.  Add  a  few  drops  of  silver  nitrate 
to  each.  (?)  (K.  134-8;  S.  187.)  Bearing  in  mind  that  salts  in  so- 
lution generally  interact  in  such  a  way  that  an  exchange  of  radicals 
takes  place,  construct  equations  for  the  actions.  How  can  you  tell 
whether  the  precipitate  is  silver  chloride  or  a  nitrate  or  both?  Which 
is  it?  (Appendix  8.)  Do  all  substances  containing  chlorine  give 
silver  chloride  by  double  decomposition  in  solution?  What  radical 
must  a  substance  contain  to  yield  silver  chloride  by  interaction  with 
silver  nitrate  solution? 

By  referring  to  the  table  of  solubilities  (Appendix  8)  try  to  predict 
which  of  the  following  solutions  will  give  precipitates  by  the  addition 
of  silver  nitrate  solution, — potassium  iodide,  potassium  bromide, 
potassium  flouride.  Give  the  reasoning  upon  which  your  prediction 
is  based.  Confirm  your  supposition  by  suitable  experiments.  (?) 

Suggest  some  other  nitrate  that  you  think  would  yield  a  precipi- 
tate of  a  chloride  from  the  solutions  of  chlorides  used  above  and  verify 
your  answer. 

b.  Filter  off  one  of  the  precipitates  of  silver  chloride  obtained  and 
add  a  few  drops  of  silver  nitrate  to  the  filtrate.  (?)  If  no  precipi- 
tate forms  what  inference  do  you  draw?  If  a  precipitate  forms  form- 


50  Laboratory  Experiments. 

ulate  your  ideas  of  how  you  would  proceed  by  this  method  to  obtain 
a  filtrate  free  from  chloride. 

c.  Add  a  few  drops  of  silver  nitrate  solution  to  solutions  of  sodium 
carbonate  and  sodium  phosphate  in  separate  test  tubes.     (?)     Add 
nitric  acid  to  the  contents  of  the  tubes  until  no  further  change  occurs. 
(?)     Add  also  a  few  drops  of  nitric  acid  to  one  of  the  pecipitates  of 
silver  chloride.       (?)     How  would  the  presence  of  a  dilute  nitric  acid 
in  solutions  of  sodium  phosphate  and  potassium  chloride  affect  the 
interaction  of  the  solutions  with  silver  nitrate? 

Add  a  few  drops  of  ammonium  hydroxide  to  silver  chloride.     (?) 
Formulate  a  test  for  a  soluble  chloride  that  will  distinguish  it  from 
a  phosphate  or  carbonate.     Would  this  test  also  distinguish  the  chlo- 
ride solution  from  the  solution  of  an  iodide  or  bromide? 
Ascertain  whether  the  tap  water  contains  chlorides.     (?) 

d.  To  a  few  drops  of  silver  nitrate  solution  add  some  ammonium 
hydroxide  and  so  obtain  a  solution  of  ammonio-silver  nitrate.     Now 
add  a  solution  of  sodium  or  potassium  chloride.     (?)     Is  the  silver 
radical   present?     Which   compounds   alone   give   silver   chloride    by 
double  decomposition? 

e.  To  solutions  of  sodium  carbonate,  sodium  phosphate,  and  so- 
dium sulphate  in  separate  test  tubes  add  a  few  drops  of  barium  chlo- 
ride solution.     Assuming  that  an  exchange  of  radicals  occurs,    write 
equations  for  the  actions,  and  determine  by  referring  to  the  table  of 
solubilities,  which  of  the  products  is  the  precipitate.     Add  a  few  drops 
of  barium  chloride  solution  to  samples  of  dilute  sulphuric  and  phos- 
phoric acids.     (?)     Add  a  little  hydrochloric  acid  to  each.     (?)     In- 
terpret the  results.     (?)     If  solutions  of  sodium  phosphate  and  so- 
dium sulphate  are  acidified  with  hydrochloric  acid,  would  you  expect 
precipitates  to  form  in  either  of  them  upon  the  addition  of  barium 
chloride?    Confirm  your  answer  by  experiments.     (?) 

Formulate  a  test  for  the  sulphate  radical.  Will  your  test  enable  you 
to  distinguish  sulphate  solutions  from  a  carbonate  solution?  From 
a  solution  of  a  phosphate? 

Suggest  an  experimental  method  of  distinguishing  between  solu- 
tions of  carbonates  and  phosphates. 

When  a  solution  of  ammonium  molybdate  is  added  to  a  solution  of 
a  phosphate,  e.  g.,  sodium  phosphate,  to  which  a  few  drops  of  nitric 
acid  have  been  added,  a  yellow  precipitate  is  obtained  either  at  once 


Acids,  Bases,  Salts.  51 


or  upon  warming  the  solution.  In  dilute  solutions  the  precipitate 
only  appears  after  the  solutions  have  stood  for  a  time.  Try  the  mo- 
lybdate  test  for  the  phosphate  radical.  (?)  The  formula  of  the  per- 
cipitate  is  complex.  Do  not  attempt  the  equation  for  the  reaction. 
(K.  491;  S.  483.) 

f.  Suppose  you  were  given  four  white  substances  and  that  you 
knew  them  to  be  a  carbonate,  chloride,  phosphate,  and  sulphate  of  a 
metal.  State  what  tests  you  would  make,  and  what  reasoning  you 
would  use,  in  order  to  identify  the  acid  radical  in  each.  Ask  the  in- 
structor for  a  specimen  of  one  of  more  of  these  salts  and  proceed  to 
identify  it.  Record  the  experiments  that  you  conduct  with  the  un- 
known and  give  the  reasoning  you  use  in  identifying  the  substance. 

33.     Examples  of  Reversible  Reactions. 

To  a  concentrated  solution  of  sodium  hydrogen  sulphate  add  about 
an  equal  volume  of  concentrated  hydrochloric  acid.  (?)  Filter  off 
the  precipitate  and  dry  it  by  pressing  it  between  filter  paper.  Dis- 
solve it  in  water  and  then  evaporate  the  solution  to  dryness  on  the 
water  bath.  Describe  and  name  the  product. 

Write  the  equation  for  the  action.  What  relation  does  this  action 
bear  to  that  in  Exp.  23d?  What  factors  determine  the  direction  of 
a  reversible  reaction?  (K.  134-8;  S.  174-9.) 

Recall  the  reduction  of  ferric  oxide  by  means  of  hydrogen  and  the 
action  of  steam  on  heated  iron.  Specify  the  conditions  that  must  be 
observed  to  complete  the  reaction  in  either  direction. 

Recall  how  Lavoisier  proved  that  oxygen  is  a  constituent  of  air. 
(K.  34.)  Write  an  equation  for  the  chemical  change.  Write  an 
equation  for  the  decomposition  of  red  precipitate  (class-room  experi- 
ment, and  also  experiment  16).  Specify  the  conditions  for  the  com- 
pletion of  the  reaction  in  either  direction. 


CHAPTER  IX. 


AMMONIA  AND  NITRIC  ACID. 

34.     Ammonia  and  Ammonium   Compounds. 

a.  Fit  a  small  flask  with  a  one-hole  stopper  and  L-tube,  and  con- 
nect the  latter  with  a  U-tube.     Put  some  water  in  the  latter,  filling 
it  to  such  height  that  air  can  return  to  the  flask  without  drawing  back 
water.     Test  the  apparatus  to  see  that   it   is  air-tight.     Prepare  a 
mixture  of  powdered  quick  lime  and  sal  ammoniac,  about   10  g.  of 
each.     Has  the  mixture  an  odor?     Place  it  in  the  flask  and  warm 
gently.     After  the  air  has  been  displaced  the  whole  of  the  gas  should 
dissolve  in  the  water.     Name  the  solution  thus  prepared. 

b.  Hold  a  glass  rod  dipped  in  concentrated  hydrochloric  acid  over 
the  solution.     (?)     Boil  about  5  cc.  of  the  solution  in  a  test  tube  and 
note  the  odor  from  time  to  time.     (?)     Compare  this  behavior  of  the 
solution  with  that  of  hydriodic  acid.     (Exp.  27.) 

c.  Neutralize  the  remainder  of  the  ammonium  hydroxide  prepared 
with  sulphuric  acid  and  evaporate  the  solution  to  dryness  on  a  water 
bath.    (?)   Scrape  the  residue  into  the  middle  of  the  dish,  invert  over  it 
over  a  small  funnel  the  stem  of  which  has  been  plugged  with  paper, 
and  heat  the  dish  strongly  for  some  time.      Is  the  sublimate  identical 
with  the  residue?     For  answer  test  it  for  the  sulphate  radical  and  the 
ammonium  radical.     The  test  for  the  latter  may  be  conducted  as  in 
a,  or  by  adding  2-3  cc.  of  sodium  hydroxide  solution  to  a  solution    of 
the  sublimate  and  then  gently  warming.     Note  the  odor.     (?) 

d.  (Two  students  working  together.)     In  the   middle  of  a  hard 
glass  tube  pack  a  plug  of  asbestos  fiber  about  0.5  cm.  thick,  and  place 
about  a  gram  of  ammonium  chloride  on  one  side  of  the  plug.      Mount 
the  tube  in  a  horizontal  position  and  place  strips  of  moist  litmus  paper 
in  both  open  ends  of  the  tube.     Heat  the  tube  steadily  at  the  point 
where  the  chloride  rests  and  observe  the  changes  in  tint  of  the  papers. 
Interpret  the  results.     (K.   18,   109;  S.  81,  73.) 

e.  Add  an  excess  of  ammonium  hydroxide  to  solutions  of  ferric 
chloride  and  aluminum   sulphate    in  separate  test  tubes.     Describe 
the   precipitates.     Besides   stating   the   color   of   the   precipitate    de- 

(52) 


Ammonia  and  Nitric  Acid.  53 

scribe  its  structure  selecting  appropriate  adjectives  from  the  follow- 
ing list: — Flaky,  flocculent,  gelatinous,  granular,  crystalline,  curdy, 
pulverulent,  finely  divided,  slimy. 

Boil  the  contents  of  the  tubes  until  the  excess  of  ammonia  is  ex- 
pelled. (?)  Then  filter  off  the  precipitates,  and  divide  each  of  the 
nitrates  into  two  portions.  Test  a  portion  of  each  for  the  ammonium 
radical.  (?)  Test  the  remaining  portions  respectively  for  the  chlo- 
ride and  sulphate  radicals.  (?) 

f.  To  solutions  of  cupric  sulphate,  zinc  sulphate,  and  barium  chlo- 
ride in  separate  test  tubes  add  ammonium  hydroxide  in  excess.     (?) 
Interpret.     (K.  443;  S.  416,  433.) 

g.  To  50  cc.  of  distilled  water  add  a  drop  or  less  of  ammonium 
chloride  solution  and  then  about  10  drops  of  Nessler's  reagent.     (?) 
Test  about  50  cc.  of  distilled  water  with  Nessler's  reagent.     (?)     Is 
the  distilled  water  ammonia  free?     (K.  157,  406.) 

h.  Observe  the  odor  of  all  the  ammonium  salts  (solids  and  liquids) 
upon  the  side-shelves.  (?)  Explain. 

Heat  about  a  gram  of  ammonium  phosphate  in  a  hard  glass  test 
tube.  Is  ammonia  one  of  the  decomposition  products?  Dissolve 
the  residue  in  water  and  test  with  litmus.  (?)  Interpret.  (S.  312; 
K.  315.) 

35.     Nitric  Acid  and  Nitrates.     Preparation  and  Properties.    (Two 

students  working  together.) 

a.  Concentrated  nitric  acid  is  very  corrosive.  Guard  against 
having  it  come  in  contact  with  the  clothes  or  skin. 

Pulverize  about  10  g.  of  sodium  nitrate  and  place  it  in  a  dry  retort. 
Add  6  cc.  of  concentrated  sulphuric  acid  and  wait  until  the  liquid  has 
moistened  the  entire  mass.  Support  the  retort  upon  a  sand  bath 
(K.  162)  and  allow  the  neck  of  the  retort  to  extend  to  the  bottom  of 
a  flask.  To  insure  complete  condensation  of  the  vapor  of  nitric  acid 
immerse  the  receiver  in  cold  water  and  keep  it  covered  with  filter 
paper  which  is  continually  moistened.  Heat  the  retort  gently  as  long 
as  the  acid  distills  off. 

What  is  the  composition  of  "concentrated"  nitric  acid?  (K.  164; 
S.  293.)  How  does  the  acid  made  differ  from  the  concentrated  acid 
in  the  laboratory?  Is  the  interaction  of  sulphuric  acid  and  sodium 
nitrate  reversible?  Assign  a  reason  for  the  fact  that  nitric  acid  is 
completely  displaced  from  the  mixture  in  the  retort. 


54  Laboratory  Experiments. 

b.  Blow  some  air  from  the  blast  through  the  acid.     (Care;  do  not 
splatter  acid  on  your  hands.)     What  is  the  color  of  pure  nitric  acid? 
To  what  substances  was  the  color  of  this  sample  due,  and  by  what 
action  was  it  formed?     Blow  the  moist  air  of  the  breath  over  this 
specimen.     (?) 

Boil  1  g.  of  sulphur  with  5  cc.  of  concentrated  nitric  acid  for  3  min- 
utes. Is  there  any  evidence  of  action?  Pour  the  clear  liquid  into 
another  test  tube,  dilute  it  with  four  times  its  volume  of  water  and 
test  for  the  sulphate  radical.  (?)  What  property  of  nitric  acid  is 
here  shown? 

Refer  to  your  record  of  the  experiment,  "Interaction  of  Metals  and 
Acids,"  and  name  the  gaseous  products  there  described  for  the  action 
of  dilute  and  concentrated  nitric  acid  with  copper  and  zinc. 

c.  In  separate  test  tubes  heat  specimens  of  lead  nitrate  and  cop- 
per   nitrate.     Apply    the    glowing    splinter    test.     (?)     What    other 
nitrates  behave  like  these?     Recall  the  heating  of  potassium  nitrate 
(Exp.  8.)     What  other  nitrates  behave  like  saltpetre?     Heat  about 
2  g.  of  ammonium  nitrate  in  a  test  tube  (?)  and  apply  the  glowing 
splinter  test.     (?)     (K.    171;   S.    300.)     How  could   you   distinguish 
between  nitrous  oxide  and  oxygen? 

d.  To  a  few  drops  of  sodium  nitrate  solution  add  5  cc.  of  a  sat- 
urated solution  of  ferrous-ammonium  sulphate.     Pour  concentrated 
sulphuric  acid  cautiously  down  the  side  of  the  test  tube  until  it  forms 
a  considerable  layer  at  the  bottom  of  the  tube.     Note  the  brown 
ring  between  the  layers  and  explain.     (K.   167;  S.  295.) 

36.  Reduction  of  Nitric  Acid  to  Ammonia.  (Optional;  Two  Stu- 
dents working  together.) 

a.  Dissolve  a  small  crystal  of  sodium  nitrate  in  2  cc.  of  water.    Add 
about  the  same  volume  of  concentrated  sodium  hydroxide  solution 
and  some  aluminum  turnings  or  chips.       Heat  gently  and  test  the 
gas  evolved  with  litmus  paper.     (?)     Note  the  odor  of  the  gas.     (?) 
Is  it  inflammable?     (K.  16,  87;  S.  68,  302.) 

b.  To  5  g.  of  zinc  in  a  test  tube  add  dilute  sulphuric  acid,  and 
when  the  evolution  of  hydrogen  is  in  progress  add  concentrated  nitric 
acid  at  the  rate  of  a  drop  a  minute.       If  ammonia  is  formed  by  com- 
plete reduction  of  nitric  acid  where  will  it  be  found,  and  in  what  con- 
dition?    Test  for  its  presence.     (?) 


CHAPTER  X. 


HYDROGEN  SULPHIDE,  SULPHUR  DIOXIDE, 
CARBON  DIOXIDE. 

37.     Hydrogen  Sulphide.    Preparation  and  Properties. 

a.  Set  up  the  apparatus  used  in  the  preparation  of  hydrogen,  but 
it  is  advisable  to  use  a  smaller  flask.  In  the  flask  is  placed  about 
20-30  g.  of  ferrous  sulphide,  and  pour  in  through  the  thistle  tube  suf- 
ficient hydrochloric  acid  to  cover  the  sulphide.  Later  if  the  action 
becomes  too  slow  add  a  little  concentrated  acid  at  intervals.  When 
the  gas  is  not  being  used  lower  the  delivery  tube  into  a  solution  of 
sodium  hydroxide.  This  prevents  its  escape  into  the  room. 

Note  the  odor  of  the  gas  and  apply  to  it  a  strip  of  test  paper, — 
filter  paper  dipped  in  lead  nitrate  solution.  (?)  When  the  air  has 
been  displaced  from  the  generator,  and  the  evolution  of  gas  is  brisk, 
collect  several  bottles  of  it  over  water.  In  connection  with  this 
operation  what  evidence,  if  any,  of  the  solubility  of  the  gas  in  water 
did  you  observe?  Is  the  gas  combustible?  Does  it  support  combus- 
tion? To  one  of  the  bottles  of  gas  covered  with  a  glass  plate  add  a 
few  drops  of  bromine  water  and  quickly  close  the  mouth  of  the  bottle 
with  the  palm  of  the  hand  and  shake.  (?)  What  evidences  of  ac- 
tion did  you  note? 

Set  fire  to  the  gas.  Describe  the  color  of  the  flame  and  note  if 
moisture  is  deposited  on  a  dry  beaker  held  for  an  instant  over  the 
flame.  Identify  by  its  odor  any  gas  formed  by  the  burning.  (?) 
Hold  a  porcelain  dish  in  the  middle  of  the  flame  for  a  few  moments. 
(?)  What  is  deposited?  Is  it  therefore  probable  that  this  substance 
exists  uncombined  in  the  interior  of  the  flame?  (S.  251.)  Pass  the 
gas  into  concentrated  sulphuric  acid.  (?)  Identify  one  of  the  pro- 
ducts of  the  interaction  by  the  odor  of  the  gas  issuing  from  the  acid. 
(?)  What  other  products  do  you  recognize?  Lead  hydrogen  sul- 
phide into  dilute  nitric  acid  and  try  to  identify  some  of  the  products. 

b.  Take  about  10-15  cc.  of  water  in  a  test  tube  and  saturate  it 
with  hydrogen  sulphide.  Test  the  solution  with  litmus  paper.  (?) 

(55) 


56  Laboratory  Experiments. 

Assign  a  name  to  the  solution.  Pour  a  drop  of  the  solution  on  the 
test  paper  used  above.  In  a  test  tube  boil  5  cc.  of  the  solution  and 
note  the  odor  of  the  vapor  from  time  to  time,  and  finally  try  to  detect 
the  presence  of  the  gas  in  the  vapor  by  means  of  the  test  paper.  Can 
the  gas  be  driven  off  completely  by  boiling?  Compare  this  behavior 
with  ammonium  hydroxide  and  hydrochloric  acid.  Expose  the  re- 
mainder of  the  solution  to  the  air  until  the  next  laboratory  period. 
Explain  the  turbidity. 

c.  Saturate  about   10  cc.  of  ammonium  hydroxide  with  the  gas. 
Test  with  litmus  paper.     (?)     To  a  few  drops  of  the  solution  just 
prepared  add  dilute  hydrochloric  acid  and  note  the  odor  of  the  es- 
caping gas.     (?)     Acidify  a  portion  of  the  caustic  soda  solution  in 
which  the  superfluous  gas  was  absorbed  in  the  course  of  the  experi- 
ments.     (?)     To  the  remainder  of  the  ammonium  hydrogen  sulphide 
solution  add  a  little  powdered  roll  sulphur,  and  shake  from  time  to 
time.     (?)     Is  sulphur  soluble  in  water?     What  is  here  to  be  inferred? 
When  the  solution  has  become  very  yellow  in   color,  filter.     What  is 
the  name  of  this  solution?     (K.  372;  S.    371.)     Acidify    the    filtrate 
with  dilute  hydrochloric  acid.     (?) 

d.  (Two  students  working  together.)     Take  six  clean  tubes  and 
obtain  2-3  cc.  of  the  solution  of  each  of  the  following  substances. 
Dilute  each  specimen  with  10-20  cc.  of  water  and  saturate  with  hy- 
drogen sulphide:     (a)  Cupric  sulphate  (?),   (b)  Lead  nitrate  (?),   (c) 
Cadmium  sulphate  (?),  (d)  Zinc  acetate  (?),  (e)  Ferrous  ammonium 
sulphate  (?),  (f)  Potassium  sulphate  (?). 

Pour  away  a  part  of  the  contents  of  each  tube,  add  a  large  excess  of 
dilute  hydrochloric  acid,  and  shake  (?).  Explain  the  results.  Divide 
the  metallic  sulphides  obtained  as  precipitates  into  two  classes,  and 
characterize  those  classes. 

Saturate  a  diluted  solution  of  ferrous  ammonium  sulphate  which  has 
previously  been  acidified  with  dilute  hydrochloric  acid  with  hydrogen 
sulphide.  (?)  Now  add  ammonium  hydroxide  drop  by  drop  until 
the  solution  shows  an  alkaline  reaction.  (?)  Filter  off  the  precipitate, 
and  lead  more  hydrogen  sulphide  into  the  filtrate.  (?)  If  no  further 
precipitate  forms,  what  may  you  infer?  If  a  precipitate  is  obtained 
how  would  you  proceed  to  obtain  a  filtrate  free  from  ferrous  sulphate? 

e.  Imagine   that   you    have    mixed   solutions   of   cupric   sulphate, 
ferrous  sulphate  and  potassium  sulphate.     State  how  you  would  pro- 


Sulphides  and  Oxides.  57 

ceed  in  the  light  of  the  facts  ascertained  in  d,  to  separate  the  metallic 
radicals  from  the  solution.  Try  your  plan,  but  before  doing  so  submit 
it  to  the  instructor  for  suggestions  and  criticism. 

f.  What  volume  of  hydrogen  sulphide  will  be  produced  by  dissolv- 
ing 20  g.  of  ferrous  sulphide,  90  per  cent  pure,  in  dilute  hydrochloric 
acid? 

g.  What  volume  of  air  is  necessary  to  burn  a  liter  of  hydrogen 
sulphide? 

h.     Suggest  a  method  of  obtaining  zinc  sulphide  from  a  solution  of 
zinc  chloride.     (?)     Try  your  method.     (?) 

38.     Sulphur  Dioxide.     Preparation  and  Properties. 

a.  Recall  the  experiments  in  which  you  identified  sulphur  dioxide. 

b.  Place  in  separate  test  tubes  specimens  of  sodium  carbonate, 
sodium  sulphite,  and  sodium  thiosulphate  (hypo),  and  add  to  each  a 
little  dilute  hydrochloric  acid.     (?)     Note  the  odor  of  the  gaseous 
products.     (?)     How  would  you  identify  the  gas  evolved  from  the 
carbonate?     (Exp.   9.)     When  the  crystals   have  dissolved   examine 
the  contents  of  the  test  tubes  and  point  out  any  difference  that  could 
be  used  to  distinguish  the  sulphite  from  the  thiosulphate. 

c.  Set  up  on  a  sand-bath  the  apparatus  used  in  preparing  hydrogen 
sulphide.     Connect   the  exit   tube  to  a   washing-bottle  of  the  type 
shown  in  Fig.  12.     This  bottle  is  to  serve  here  as  a  safety  trap  (how 

and  why?)  and  is  to  be  left  empty.  The 
shorter  tube  of  the  washing  bottle 
should  be  toward  the  generator.  Why 
not  the  reverse?  Into  the  flask  put 
10-15  g.  of  copper  chips  or  turnings, 
and  through  the  thistle  tube  pour  about 
25  cc.  of  concentrated  sulphuric  acid. 
Test  the  apparatus  to  see  that  it  is  air- 
tight. Heat  the  flask  and  contents 
**£•  12.  until  a  gas  begins  to  evolve.  Then 

moderate  the  heating  so  that  a  slow  steady  stream  of  gas  will  escape. 

While  the  gas  is  not  being  collected  pass  it  into  water  and  a  solution 

of  sodium  hydroxide.     Save  these  solutions  as  well  as  the  contents  of 

the  generator. 


1 


58  Laboratory  Experiments. 

Collect  four  bottles  of  gas  by  displacement  of  air.  Invert  the  last 
bottle  of  gas  collected  over  water.  (?)  If  any  gas  remains  in  the 
bottle,  what  should  you  expect  it  to  be? 

Test  the  odor  of  the  gas,  and  test  its  action  on  moist  litmus  paper. 
(?) 

In  the  third  bottle  place  the  following  substances: — Litmus  paper, 
colored  cloth,  paper  on  which  is  ordinary  writing  and  printing,  green 
grass  or  a  flower.  Stopper  the  bottle  and  allow  it  to  stand  until  the 
next  laboratory  period.  (?)  Then  explain  the  changes. 

Does  the  gas  burn?     Does  it  support  combustion. 

c.  Test  the  aqueous  solution  of  the  gas  with  litmus.     (?)     Name 
it.     Boil  about  5  cc.  of  the  solution  in  a  test  tube  and  note  the  odor 
from  time  to  time.     (?)     Finally  test  the  remaining  liquid  with  litmus 
paper.     (?)     What  other  solutions  of  gases  examined  before  behaved 
in  the  same  way?     Test  a  portion  of  the  aqueous  solution  for  the 
sulphate  radical  by  adding  barium  chloride  and  then  pure  hydrochloric 
acid.     (?)     Is  the  action  reversible?     Add  bromine  water  to  another 
portion  of  the  sulphurous  acid  until  the  tint  is  permanent,  and  then 
test  for  the  sulphate  radical.     (?)     What  property  of  bromine  water  is 
shown  by  this  experiment.     (?) 

Add  potassium  permanganate  solution  to  another  portion  of  the 
solution  a  few  drops  at  a  time,  as  long  as  the  pink  color  disappears. 
Test  the  solution  for  the  sulphate  radical. 

Leave  the  remainder  of  the  aqueous  solution  in  the  bottle  for  several 
days  exposed  to  the  air.  Then  test  for  the  sulphate  radical.  (?) 

d.  Place  about  10  cc.  of  potassium  permanganate  solution  in  each 
of  three  test  tubes.     In  three  more  place  solutions  of  potassium  di- 
chromate.     Acidify  each  solution  with  2  cc.  of  concentrated  sulphuric 
acid.     Into  one  of  the  samples  of  permanganate  solution  pass  sulphur 
dioxide  or  add  the  aqueous  solution  of  the  gas  until  the  pink  color  is 
completely  discharged.     Treat  one  of  the  dichromate  solutions  in  the 
same  way.     (?)     Into  another  pair  of  the  solutions  lead  hydrogen 
sulphide  until  no  further  change  occurs.     (?)     In  the  third  set  place 
a  few  pieces  of  granulated  zine.     (?)     Explain  the  actions.     (K.  500, 
490;    S.    492,    278,   253.)     What    property    of    sulphurous    acid    and 
hydrogen  sulphide  is  illustrated  by  these  experiments?   Classify  all  the 
reactions  of  this  section  under  one  head. 


Sulphides  and  Oxides.  59 

e.  To  a  small  portion  of  the  sodium  hydroxide  used  to  absorb  the 
superfluous  gas,  add  dilute  hydrochloric  acid  until  the  solution  reacts 
acid  toward  litmus.     (?)     If  any  gas  is  evolved  note  its  odor.     (?) 
To  another  small  portion  of  the  solution  add  a  little  bromine  water. 
Test  both  solutions  for  the  presence  of  the  sulphate  radical.     (?) 

f.  Describe  the  contents  of  the  generator.     If  any  copper  chips 
remain  remove  them  from  the  residue  after  diluting  with  a  small  volume 
of  water,  rinse  them  and  return  to  the  shelf  bottle.     Formulate  a  plan 
of  obtaining  crystals  of  blue  vitriol  from  the  residue.     (?)     Submit 
your  plan  to  the  instructor  for  criticism  and  finally  show  him  your 
yield  of  crystals. 

39.  Carbon  Dioxide.  Preparation  and  Properties.  (Two  students 
working  together.) 

a.  Enumerate    the    reactions   of    previous   experiments   in    which 
carbon  dioxide  was  identified.     What  property  of  carbon  dioxide  have 
you  used  for  its  identification? 

b.  Place  a  few  small  pieces  of  marble  or  limestone  in  hard  glass 
test  tube  fitted  with  a  cork  and  L-tube,  and  heat  strongly.     Pass  the 
gas  through  5  cc.  of  lime  water  in  a  test  tube.     (?)     What  is  the  residue? 
Place  it  in  an  equal  bulk  of  water,  and  set  aside  for  a  few  minutes. 
Then  add  about  20  cc.  of  water  and  shake  vigorously  and  filter.     Test 
the  filtrate  with  litmus  paper.     (?)     Name  the  filtrate. 

c.  Set  up  a  generating  flask  of  the  type  used  in  preparing  hydrogen , 
and  connect  it  with  a  wash-bottle  containing  some  distilled  water. 
Put  in  the  flask  pieces  of  marble  and  pour  upon  them  dilute  hydro- 
chloric acid.     Collect  the  gas  in  bottles  by  downward  displacement 
of  air  or  over  water.     What  is  the  function  of  the  wash-bottle?     Verify 
your   answer   experimentally.     (?)     What    substance   could   be   sub- 
stituted for  marble  in  this  experiment? 

To  one  bottle  of  the  gas  add  some  distilled  water,  close  with  the  hand, 
and  shake.  Test  the  solution  with  litmus.  (?) 

Has  the  gas  odor?  Does  it  support  combustion?  Pour  a  bottle  of 
the  gas  over  a  lighted  taper.  (?) 

Collect  a  bottle  of  the  gas  over  water  and  keep  it  inverted  over  water 
until  the  next  laboratory  period.  (?) 

What  name  is  sometimes  applied  to  the  aqueous  solution  of  the  gas? 
Boil  an  aqueous  solution  of  the  gas  for  a  minute  and  then  test  the  re- 
maining liquid  with  lime  water.  (?) 


60  Laboratory  Experiments. 


Pass  the  superfluous  gas  from  the  generator  into  50  cc.  of  ammonium 
hydroxide  and  store  it  in  a  corked  vessel  for  future  use. 

By  means  of  a  tube  blow  air  from  the  lungs  through  lime-water.  (?) 
How  could  you  determine  the  proportion  of  carbon  dioxide  in  the  air 
expelled  from  the  lungs? 

Explain  the  statement, — The  carbon  dioxide  present  in  a  sample 
of  water  was  fixed  by  adding  a  few  cc.  of  strong  caustic  potash. 

What  is  soda?  Soda  ash?  Washing  soda?  Baking  soda?  Caustic 
soda?  What  is  soda  water,  and  how  did  it  get  its  name?  How  can 
you  prove  that  plain  soda  water  contains  no  soda  in  solution?  How 
is  soda  water  made  commercially?  What  causes  Bromo  Seltzer, 
Sedlitz  powders  and  the  like  to  effervesce  when  water  is  added?  What 
is  baking  powder  made  of?  Identify  one  of  the  products  obtained 
when  baking  powder  is  moistened  with  water?  How  is  the  rising  of 
dough  and  bread  explained? 

d.  In  a  flask  place  about  100  cc.  of  lime  water  and  pass  a  stream 
of  carbon  dioxide  persistently  through  the  solution.  (?)  Divide  the 
clear  liquid  into  three  parts.  (If  not  quite  clear,  filter  it).  Boil  one 
portion  gently  in  a  small  flask  fitted  with  a  delivery  tube  and  pass  the 
vapors  into  lime  water.  (?)  What  is  the  solid  in  the  flask?  Explain 
temporary  hardness  as  applied  to  waters.  To  the  second  portion  add 
lime  water.  (?) 

To  the  remainder  of  the  temporary  hard  water  add  1  cc.  of  soap 
solution  and  shake.  Does  a  lather  remain.  Is  a  precipitate  present 
in  the  liquid?  Save  the  sample. 

Add  a  few  drops  of  soap  solution  to  approximately  the  same  volume 
of  distilled  water,  and  shake.  Does  a  lather  form?  If  not,  continue 
adding  soap  solution,  a  few  drops  at  a  time,  until  a  lather  remains  on 
the  liquid  for  a  minute.  Determine  the  volume  of  soap  solution  re- 
quired to  form  a  lather  in  the  temporary  hard  water.  What  is  the 
precipitate  that  forms  in  the  temporary  hard  water? 

Summarize  the  methods  of  softening  temporary  hard  water. 

Dissolve  a  small  crystal  of  magnesium  sulphate  in  a  test  tube  full 
of  distilled  water.  To  another  sample  of  water  add  a  pinch  of  pow- 
dered gypsum;  shake  persistently,  and  then  filter.  You  now  have 
two  samples  of  permanent  hard  water.  Add  a  few  drops  of  soap 
solution  to  each  and  determine  the  volume  of  soap  solution  required 


Sulphides  and  Oxides.  61 

to  form  a  lather.  What  is  the  precipitate  that  forms?  (K.  253; 
S.  335.) 

Give  a  method  of  softening  permanent  hard  water  that  does  not 
involve  the  use  of  soap.  (S.  392.) 

What  is  the  commercial  name  of  crystallized  magnesium  sulphate? 

40.  Sodium   Carbonate   by  the   Solvay   Process.     (Two  students 
working  together.     Optional.) 

Unite  the  solutions  that  were  obtained  by  passing  the  waste  carbon 
dioxide  of  the  previous  experiment  into  ammonium  hydroxide,  and 
saturate  the  solution  with  sodium  chloride  by  persistently  shaking 
in  a  corked  flask.  Decant  the  clear  liquid  into  another  bottle,  fitted 
with  two  tubes,  one  of  which  reaches  to  the  bottom.  If  because  of 
delay  a  precipitate  has  appeared,  proceed  without  decanting  the 
solution.  Through  the  longer  tube  pass  in  carbon  dioxide  until  the 
solution  is  saturated.  This  operation  will  require  at  least  half  an 
hour,  and  generally  longer.  During  the  absorption  of  carbon  diox- 
ide the  exit  tube  should  be  closed  to  prevent  waste  of  the  gas.  Close 
the  tubes  with  caps  of  rubber  tubing  plugged  with  glass  rods  and  set 
aside  over  night.  (?)  Filter  off  the  deposit  and  dry  by  pressing 
between  filter  papers. 

Dissolve  in  water  a  little  of  the  solid,  and  test  the  reaction  of  the 
solution  with  litmus.  (?)  If  the  solution  is  not  acid  explain  why  it 
is  not  so. 

Add  any  dilute  mineral  acid  to  a  part  of  the  solid.     (?) 

Heat  the  rest  in  a  test  tube  fitted  with  a  cork  and  delivery  tube, 
and  pass  the  gas  into  lime  water.  When  gas  ceases  to  be  given  off, 
dissolve  the  cold  residue  in  a  very  little  water,  test  the  reaction  of  the 
solution  with  litmus  paper,  and  set  it  aside  to  crystallize  in  an  open, 
shallow  dish.  (?)  Explain  the  reaction  with  litmus.  (?)  Dry  the 
crystals  and  try  the  effect  of  acid  upon  them.  (?) 

What  are  the  commercial  names  of  sodium  hydrogen  carbonate 
and  crystallized  sodium  carbonate? 

41.  Identification  of  Unknown  Salts. 

a.  Make  a  list  of  the  negative  radicals  studied  in  the  laboratory 
experiments,  and  place  opposite  each  the  gases,  if  any,  that  would 
be  evolved  by  the  addition  (a)  of  dilute  hydrochloric  acid,  and  (b) 


62  Laboratory  Experiments. 

a  little  concentrated  sulphuric  acid  and  warming,  to  soluble  salts  of 
the  radicals.     Arrange  your  data  in  tabular  form. 

b.  Apply  to  the  instructor  for  two  unknown  substances,  and  as- 
certain by  means  of  experiments  the  acid  radical  in  each.  Report 
your  conclusion  to  the  instructor  with  a  report  giving  the  reasoning 
used  in  identifying  the  unknown. 


CHAPTER  XI. 


THE  ATMOSPHERE,  FLAME,  OXIDATION  AND 
REDUCTION. 

42.     Components  of  Air.     (Two  students  working  together.) 

a.  Place  3-5  cc.  of  clear  barium  hydroxide  solution  in  the  bottom 
of  a  small  beaker  and  leave  it  exposed  to  the  air  until  the  end  of  the 
laboratory  period.     (?)     Relate  this  experiment  to  the  observations 
made  when  air  from  the  lungs  was  blown  through  lime  water.     (Exp. 
39  c.) 

b.  (Quant.)    On  watch  glasses  expose  to  the  air  weighed  specimens 
of  blue  vitriol,  fused  calcium  chloride,  Glauber's  salt,  and  metallic 
sodium.      Toward   the   close   of    the  laboratory  period  weigh  again, 
and  make  another  set  of  weigh-  ings  during  the  following  laboratory 
period.     Construct    a    table    showing  the   changes  in  weight  of  the 
substances.     Account  for  the  observed  variations  in  weight  and  write 
equations  indicating  the  changes  that  have  occurred  in  the  different 
substances. 

Note  9.  When  you  are  requested  to  give  an  interpretation  of  re- 
sults or  furnish  an  explanation,  please  bear  in  mind  that  an  equation 
alone  is  not  a  sufficient  answer. 

Add  a  few  drops  of  acid  to  the  material  that  resulted  from  the  ex- 
posure of  sodium  to  the  air.  (?) 

Explain  the  terms,  deliquescent,  efflorescent,  hygroscopic. 

c.  Place  about  5  g.  of  iron  filings,  moistened  with  a  few  drops  of 
water  in  the  bowl  of  a  retort.     Clamp  it  vertically  on  a  retort  stand 
so  that  the  neck  dips  about  6  cm.  below  the  surface  of  some  water 
in  a  beaker  of  about  300  cc.  capacity.     With  a  file  mark  where  the 
water  stands  in  the  neck  of  the  retort.     Put  the  stopper  in  place  and 
cover  it  with  a  little  melted  paraffine  to  insure  the  apparatus  being 
air-tight.     Set  aside  until  next  laboratory  period,  but  mark  the  posi- 
tion of  the  water  in  the  neck  of  the  retort  at  the  close  of  the  labora- 
tory period.     Read  the  barometer  and  note  the  room  temperature. 
Upon  resuming  the  experiment  mark  the  position  of  the  water  with 

(63) 


'64  Laboratory  Experiments. 


a  file.  Read  the  barometer,  and  take  the  room  temperature.  Insert 
a  cork  in  the  end  of  the  neck  of  the  retort  in  order  to  prevent  the  water 
from  running  out.  Test  the  gas  left  in  the  bowl  by  means  of  a  burn- 
ing splinter  of  wood.  Is  it  air?  Why  is  the  volume  different  from 
the  original  air?  To  what  extent  do  the  changes  in  pressure  and 
temperature  account  for  the  observed  changes  in  volume?  Has  the 
iron  changed  in  any  way?  What  does  the  contraction  represent? 
Measure  the  volume  which  represents  the  contraction  and  also  that 
of  the  gas  that  remains.  Find  the  percentage  of  each  of  the  original 
volume. 

The  experiment  may  be  conducted  with  a  glass  cylinder  in  place 
of  the  retort. 

d.  (Quant.)  A  measured  volume  of  air  is  driven  over  heated 
copper,  and  the  residual  gas  collected  over  water  and  measured. 
Place  several  wads  of  copper  turnings  in  a  hard  glass  tube  20  cms. 
long.  Fit  the  tube  with  one-hole  stoppers  provided  with  glass  tubes. 
Connect  one  of  the  tubes  to  the  aspirator  used  in  determining  the 
weight  of  a  liter  of  oxygen.  Connect  the  other  to  a  bottle  containing 
a  measured  volume  of  air.  The  air  container  should  be  a  bottle  of  at 
least  500  cc.  capacity  fitted  with  a  two-hole  rubber  stopper  carrying 
a  delivery  tube  and  a  drop  funnel.  When  the  hard  glass  tube  has 
been  heated  to  redness  start  the  flow  of  air  over  the  copper,  and 
maintain  a  slow  steady  current.  It  is  well  to  screen  the  aspirator 
from  the  heat  of  the  burners.  Ascertain  the  volume  of  the  redisual 
air  from  the  volume  of  water  displaced  from  the  aspirator  making 
proper  adjustment  of  water  levels.  Find  the  percentage  of  oxygen 
by  volume.  What  are  the  principal  constituents  of  the  residual  gas? 
Does  it  still  support  combustion? 

Examine  the  turnings,  and  by  bending,  twisting,  and  scraping  re- 
move most  of  the  scale.  Compare  it  with  the  oxides  of  copper  on 
the  side-shelf.  (?)  Place  the  scale  in  the  bottom  of  a  test  tube, 
cover  it  with  a  thin  layer  of  powdered  charcoal,  and  heat  gently. 
When  cold  pour  the  contents  of  the  tube  into  a  small  beaker  and 
separate  the  components  of  the  mixture  by  letting  a  small  stream  of 
water  flow  into  the  beaker.  What  is  the  residue  in  the  beaker?  How 
was  copper  oxide  reduced  in  an  earlier  experiment? 


Air,  Oxidation,  Reduction.  65 

43.  Bunsen  Flame. 

a.  Prepare  a  clean,  bright  copper  wire  and  explore  the  non-lumi- 
nous flame  with  it.     Watch  the  play  of  colors  on  the  wire  and  ex- 
plain.    Cut  off  the  gas  supply  until  the  flame  is  about  6  cm.  high, 
and  close  the  air  holes  until  a  luminous  point  appears  at  the  apex  of 
the  inner  cone.     Heat  portions  of  the  wire  coated  with  oxide  in  the 
luminous  region  of  this  flame.     Before  withdrawing  the  wire  lower 
it  into  the  inner  cone  of  unburnt  gas  to  cool.     (?) 

b.  Make  a  borax  bead  by  dipping  the  glowing  end  of  the  platinum 
wire  in  a  bit  of  borax  on  a  watch  glass,  and  then  heating  in  the  flame. 
Observe  the  behavior  of  the  borax  and  explain.     (K.  302;  S.  350.) 
Bring  the  hot  borax  head  in  contact  with  a  minute  particle  of  cupric 
oxide  and  heat  it  in  the  oxidizing  region  of  the  flame  until  the  particle 
has  dissolved.     What  is  the  color  of  the  bead  when  cold?     Now  ad- 
just the  flame  for  reducing  and  heat  the  bead  persistently  in  the 
luminous  apex.     (?)     Explain.     Finally  reheat  the  bead  in  the  oxi- 
dizing flame.     (?) 

Remove  the  head  from  the  wire  by  fusing  it  again,  and  giving  the 
wire  a  short,  sharp  jerk,  or  by  plunging  the  molten  bead  into  a  little 
dilute  hydrochloric  acid.  Prepare  a  fresh  bead,  (it  should  be  clear  and 
transparent,  when  cool)  and  disslove  in  it  a  speck  of  manganese  dioxide. 
If,  after  heating  in  the  oxidizing  flame  the  bead  is  black,  too  much  of 
the  dioxide  has  been  used.  Throw  the  molten  bead  off  and  try  again. 
Heat  in  the  reducing  flame  and  finally  again  in  the  oxidizing  flame. 
(?)  Explain. 

Try  the  behavior  of  a  cobalt  compound  with  borax  in  the  oxidizing 
and  reducing  flames.  (?) 

What  is  the  source  of  the  oxygen  in  these  experiments? 

c.  Pulverize  a  few  particles  of  potassium  sulphate.     Adjust  the 
Bunsen  flame  for  reducing  purposes.     Heat  the  platinum  wire,  touch 
the  salt  with  it,  and  heat  the  adhering  powder  steadily  in  the  luminous 
apex.     After  a  minute  or  two  withdraw  the  bead,  place  it  upon  a  clean 
silver  coin  and  moisten  with  a  drop  of  water.     (?)     Explain  the  result. 
(S.  420.)     What  are  the  reducing  agents  here. 

44.  Destructive  Distillation. 

a.  Arrange  the  apparatus  shown  in  Fig.  12.  Fill  the  hard  glass 
test  tube  about  half  full  of  sawdust.  Heat  gently  at  first  and  then 

Chem. — 5 


Laboratory  Experiments. 


strongly  until  no  further  change  occurs.  (?)  While  heating  bring 
a  flame  to  the  end  of  the  jet.  (?)  Note  the  odor  of  the  gas,  and  its 
reaction  toward  moist  litmus  paper.  What  is  the  residue?  Reserve 
it.  Test  the  liquid  in  the  bottle  with  litmus  paper.  (?) 

b.  Clean  the  apparatus  or  replace  it  with  a  fresh  set  and  repeat 
the  experiment  using  bituminous  coal.     (?)     Save  the  residue. 

c.  Repeat  the  experiment  using  gelatine.     (?) 

d.  Recall  the  heating  of  cane  sugar.     (Exps.  8,  13  a.)     Was  the 
experiment  an  example  of  destructive  distillation? 

f.     (Quant.)     (Two  students  working  together.)     Weigh  a  covered 


Fig.  13. 

porcelain  crucible,  and  place  in  it  about  2  g.  of  powdered  coal,  and  weigh 
again.  Place  the  crucible  on  a  pipe  stem  triangle  and  heat  it  strongly 
in  a  flame  about  20  cm.  high.  Make  sure  that  the  lid  fits  closely  on 
the  rim  of  the  crucible.  Continue  the  heating  for  five  minutes  after 
the  gases  no  longer  burn  between  the  crucible  and  lid.  Cool,  weigh, 
and  from  the  loss  in  weight  compute  the  percentage  of  volatile  matter. 
Weigh  out  a  second  sample  of  the  powdered  coal,  about  1  g.,  and 
heat  it  with  the  cover  on  until  the  volatile  matter  is  gone.  Then 
remove  the  lid  and  heat  strongly,  inclining  the  crucible  slightly  until 
constant  weight  is  obtained.  From  the  data  now  in  hand  compute 
the  percentage  of  fixed  carbon  and  ash. 


Air,  Oxidation,  Reduction.  67 

Suggest  a  method  of  estimating  the  moisture  in  a  sample  of  coal. 

Did  you  note  any  evidence  that  would  show  that  sulphur  compounds 
are  present  in  the  sample  of  coal. examined.  What  is  meant  by  the 
"volatile  combustible"  matter  of  coal? 

Compare  your  percentages  with  the  values  obtained  by  two  other 
students  who  made  determinations  on  the  same  sample  of  coal,  and 
discuss  jointly  and  with  the  instructor  the  factors  responsible  for  the 
differences  in  value.  (?) 

45.     Reductions  and  Oxidations.     (Two  students  working  together.) 

a.  Mix  intimately  in  the  mortar  5  g.  of  finely  pulverized  potassium 
nitrate  with  2  g.  of  the  charcoal  or  coke  obtained  in  the  previous 
experiment.     Drop  portions  of  the  mixture  in  a  red-hot  crucible.     (?) 
What  gases  are  evolved?     Test  the  residue  with  an  acid.     (?) 

b.  Mix  on  paper  intimately   (care)   5  g.  of  powdered  potassium 
nitrate  and  2  g.  of  flowers  of  sulphur.     Throw  the  mixture  in  small 
portions  into  a  red-hot  crucible.     (?)     What  gases  are  evolved?     Dis- 
solve the  residue  in  distilled  water  and  add  a  few  cc.  of  pure  dilute 
hydrochloric  acid  and  then  barium  chloride  solution.     (?) 

c.  Make  some  "touch  paper",  or   write  on  porous  paper,  with  a 
solution  of  potassium  nitrate.     When  it  is  dry,  apply  a  lighted  match. 
(?)     Make  a  "slow  match"  or  fuse  by  placing  a  piece  of  narrow  tape 
for  a  few  minutes  in  a  boiling  solution  of  lead  nitrate  (10  cc.).     When 
dry  light  it  and  hold  some    over   white  paper  while    burning.      (?) 
What  is  "Greek  Fire"?     What  explosives  are  used  in  fireworks? 

Write  the  equation  for  the  explosion  of  gunpowder  (S.  366;  K.  349.) 
What  product  is  the  white  smoke? 

d.  Mix  a  pinch  of  powdered  arsenic  trioxide  with  pulverized  char- 
coal prepared  in  the  previous  experiment.     Heat  the  mixture  strongly 
in  a  dry  test  tube.     (?) 

e.  (Quant.)     Weigh  on  a  watch  glass  about  2  g.  of  dry  powdered 
sodium  nitrate.     Transfer  the  salt  to  a  piece  of  smooth  paper  and  mix 
it  thoroughly  with  about  seven  times  its  bulk  of  iron  powder.     In- 
troduce the  mixture  into  a  test  tube,  and  arrange  for  the  collection 
of  gas  by  means  of  the  aspirator  used  in  determining  the  weight  of  a 
liter  of  oxygen.     Heat,  beginning  at  the  top  of  the  mixture,  till  no  more 
gas  is  delivered.     Ascertain  the  volume  of  the  nitrogen  from  the  weight 
of  water  displaced.     Reduce  the  volume  of  the  gas  to  standard  con- 


68  Laboratory  Experiments. 

ditions,  and  calculate  its  weight  assuming  that  its  density  is  0.001273. 
What  percentage  of  nitrogen  is  present  in  sodium  nitrate?  What 
percentage  is  represented  by  the  formula?  Instead  of  sodium  nitrate, 
potassium  nitrate  or  barium  nitrate  may  be  used.  How  is  sodium 
nitrate  prepared? 

f.  (Optional.)     Mix   thoroughly   equal   bulks   of    finely   powdered 
marble  and  magnesium  powder.     Fill  a  test  tube  to  a  depth  of  5  cm. 
with  the  mixture,  clamp  in  an  inclined  position,  and  heat  the  top  layer 
in  the  Bunsen  flame  until  the  reaction  begins.     Apply  a  light  to  the 
gases  escaping  while  the  reaction  is  in  progress.     (?)     If  the  reaction 
does  not  go  to  completion  by  itself,  heat  the  contents  of  the  tube  again, 
but  be  careful  to  point  the  tube  away  from  the  face  during  the  heating. 
Allow  the  test  tube  to  cool,  add  a  little  water,  and  transfer  the  material 
to  a  small  beaker  and  boil.     If  any  gas  is  liberated  upon  the  addition 
of  water,  note  its  odor  and  see  whether  it  burns.     Acidify  the  mixture 
with  hydrochloric  acid  (?)  and  when  all  action  has  ceased,  filter  and 
wash  the  residue  with  water.     Dry  it  on  the  steam-bath  and  prove 
that  it  is  carbon.     (?)     What  is  the  reducing  agent  in    this    action? 
Why  was  the  hydrochloric  acid  added?     Consider  carefully  the  original 
reaction  with  a  view  of  establishing  what  gas  burned  at  the  mouth  of 
the  tube,  as  the  glow  spread  through  the  mixture.     (?) 

Evaporate  the  filtrate  t,o  dryness  on  a  sand-bath  in  an  evaporating 
dish,  and  finally  heat  whilst  stirring  over  the  free  flame.  Expose  the 
product  to  the  air  until  the  following  laboratory  period.  (?)  What 
compounds  previously  examined  behaved  similarly.  Interpret  the 
result. 

g.  Preparation  of  Amorphous  Silicon.     (Optional.) 

Prepare  an  intimate  mixture  of  clean,  finely  powdered  sand  and 
magnesium  powder,  using  equal  weights  of  the  substances,  and  about 
5  g.  of  each.  Fill  a  test  tube  to  a  depth  of  6  cm.  with  the  mixture  and 
heat  in  the  Bunsen  flame  as  directed  in  the  previous  experiment.  (?) 
Digest  the  residue  with  dilute  hydrochloric  acid,  and  apply  a  lighted 
match  to  the  gaseous  products.  (?)  Is  there  any  evidence  that 
silicon  hydride  is  present  in  the  escaping  gas?  After  the  action  has 
ceased,  filter  and  wash  the  residue  with  water.  Dry  it  by  warming, 
and  divide  into  two  parts.  Heat  one  portion  strongly  on  an  inverted 
crucible  cover.  (?)  Place  the  other  portion  in  a  test  tube  and  digest 
it  with  caustic  potash  solution.  (?) 


Air,  Oxidation,  Reduction.  69 

46.  Alcohol  by  Fermentation.  (Optional;  two  students  working 
together.) 

Dissolve  20  g.  of  glucose  syrup  or  molasses  in  ISO  cc.  of  water.  Break 
up  a  half  yeast  cake  in  luke  warm  water  until  it  is  thoroughly  mashed. 
Add  the  yeast  to  the  syrup  solution,  fill  a  flask  to  the  base  of  the  neck 
with  the  mixture,  and  fix  a  plug  of  cotton  in  its  neck.  Set  the  whole 
aside  in  a  warm  place  for  3-4  days.  Warm  the  solution  and  test  the 
gas  that  is  given  off  for  carbon  dioxide.  (?) 

Filter  the  liquid  and  place  it  in  a  larger  flask,  fitted  with  a  cork  and 
exit  tube  connected  to  a  condenser  (K.  37;  S.  26.)  Distill  off  about 
50  cc.,  and  note  the  odor  of  the  distillate.  (?)  Place  the  distillate 
in  a  distilling  flask  fitted  with  a  thermometer,  (Exp.  26),  boil  with  a 
small  flame  and  catch  the  fraction  that  passes  over  between  75-93°. 

Again  note  the  odor  of  the  distillate.  (?)  Test  it  with  litmus  paper. 
(?)  Use  a  few  drops  to  see  whether  it  will  burn.  To  the  rest  add  a 
crystal  of  iodine  and  just  enough  sodium  hydroxide  solution  to  dissolve 
it.  Heat  gently  and  then  cool.  (?)  The  yellow  crystalline  substance 
that  forms  is  iodoform,  and  this  is  the  iodoform  test  for  alcohol.  Note 
the  odor  of  the  iodoform. 


CHAPTER  XH. 
IONIZATION. 


V'/> 

fflJL 


47.  Ionic  Materials.  (Two  students  working  together.) 
a.  Secure  from  the  store-room  an  electrolytic  cell  of  the  type 
shown  in  Fig.  14.  Set  it  up  on  the 
table  equipped  for  electrolytic  ex- 
i,  by  connecting  in  series 
between  the  terminals  of  a  110  volt 
direct  current,  an  8  or  16  candle 
power  incandescent  lamp  and  the 
cell  wires.  If  the  circuit  is  not 
provided  with  simple  switch,  a 
single  pole,  single  throw  knife 
switch  should  be  placed  in  the 
circuit.  Draw  a  diagram  of  the 
connections  as  you  have  made  them. 
Bring  from  the  side-shelf  in  test 
tubes  the  following  materials  to  be 
tested  qualitatively  for  electrical 
conductivity: 

Dilute  hydrochloric  acid, 
Concentrated  hydrochloric  acid. 
Distilled  water, 


Fig.  14. 
Tap  water, 
Dilute  acetic  acid, 

Dilute  potassium  hydroxide  solution, 
Dilute  ammonium  hydroxide, 
Cane  sugar  solution, 
Cupric  sulphate  solution, 
Sodium  chloride  solution, 
Sodium  chloride  (dry  salt), 
Kerosene. 

(70) 


lonization.  71 

Remove  from  the  cell  the  stopper  carrying  the  platinum  electrodes 
and  introduce  sufficient  of  the  dilute  hydrochloric  acid  to  have  the 
electrodes  covered  with  liquid.  Exercise  care  in  handling  the  elec- 
trodes. Have  them  parallel  at  the  outset  and  then  do  not  shift  or 
bend  the  plates  during  the  measurements.  Replace  the  electrodes 
and  make  sure  that  the  wires  leading  into  the  glass  tubes  make  contact 
with  the  mercury.  Now  close  the  switch  for  a  moment  and  observe 
any  action  within  the  cell,  and  note  the  brightness  of  the  lamp.  (?) 
Empty  the  cell,  rinse  it  and  the  electrodes  well  with  water,  and  then 
continue  the  examination  of  the  other  materials.  (?)  Dry  the  cell 
before  introducing  the  dry  sodium  chloride. 

Arrange  the  substances  in  the  order  of  their  conductivity. 

b.  Perform  an  experiment  that  will  throw  some  light  on  how  the 
concentration  of  the  dissolved  substance  influences  the  resistance  of  an 
electrolyte. 

Taking  the  above  substances  as  typical  of  the  classes  to  which  they 
belong,  indicate  those  classes  of  compounds  which  readily  conduct,  and 
are  therefore  ionized  in  aqueous  solution. 

c.  In  aqueous  solutions  the  radicals  hitherto  considered  are  the 
ions.     (K.  428;  S.  219.)     Give  the  names  and  symbols  of  the  three 
chief  constituents  in  addition  to  water  in  aqueous  solutions  of  sodium 
chloride,  sodium  hydroxide,  cupric  sulphate,  and  sulphuric  acid. 

Write  in  ionic  form  the  equations  that  appear  in  your  notebook  for 
the  experiments  in  the  Chapter  on  Acids,  Bases,  and  Salts. 

What  evidence  can  you  cite  in  favor  of  the  view  that  the  ions  are 
electrically  charged  particles? 

Define  acid  and  base  in  terms  of  the  hypothesis  of  ions. 

Formulate  the  following  actions  in  terms  of  the  ionic  hypothesis: 

The  displacement  of  hydrogen  from  dilute  acids  by  zinc. 

The  displacement  of  iodine  from  solutions  of  iodides  by  chlorine. 

The  displacement  of  sulphur  from  hydrogen  sulphide  by  iodine. 

The  precipitation  of  cupric  sulphide  by  the  reaction  of  hydrogen 
sulphide  with  a  solution  of  cupric  sulphate. 

The  evaporation  of  water  from  a  salt  solution. 


CHAPTER   XIII. 


ELECTIVE  QUALITATIVE  AND  QUANTITATIVE  EXPERI- 
MENTS. 

48.  Determination    of    Carbon    Monoxide    from    a    Carbonate. 
(Quant.)     (Two    students    working    together.)     Weigh    on    a    watch 
glass  about  1.5  g.  of  finely  powdered  Iceland  spar,  or  dried  precipitated 
chalk,  and  proceed  as  in  the  determination  of  nitrogen  from  a  nitrate, 
Exp.  45  e,  but  use  zinc  dust  instead  of  iron,  and  a  hard  glass  tube 
instead  of  an  ordinary  test  tube.     Continue  the   heating  until   gas 
ceases  to  be  given  off.     Find  the  volume  of  the  gas  under  standard 
conditions,  and  calculate  its  weight  assuming  its  density  to  be  0.0012499 
Calculate  the  percentage  of  carbon  monoxide  in  terms  of  the  carbonate, 
and  compare  it  with  the  value  that  may  be  deduced  from  the  formula 
of  the  carbonate. 

Apply  a  lighted  match  to  the  gas.  (?)  What  are  some  of  its  other 
chemical  properties?  What  is  its  physiological  action?  Suggest  a 
way  of  demonstrating  its  presence  in  tobacco  smoke. 

49.  Reduction    of    Sulphide.     (Quant.)     (Two    students    working 
together.)     Prepare  approximately   2   g.   of  cupric  sulphide   starting 
with  blue  vitriol,  CuSO4.5H2O.     Calculate  the  quantity  of  blue  vitriol 
that  will  be  required,  dissolve  it  in  water,  and  pass  hydrogen  sulphide 
into  the  solution  until  the  precipitation  is  completed.     To  ascertain 
when  the  solution  is  saturated  with  hydrogen  sulphide,   filter  off   1 
to   2   cc.   and  lead  hydrogen  sulphide  into  the  filtrate.     (?)     When 
precipitation  is  complete,  allow  the  precipitate  to  settle,  and  then 
wash  it  thoroughly  by  decantation  with  water  containing  hydrogen 
sulphide.     (?)     Then    filter   and    wash    once    with    water   containing 
hydrogen  sulphide.     What  ion  is  being  removed  in  washing?     Test 
for  its  presence  in  the  last  filtrate.     Dry  the  cupric  sulphide  on  a  water 
bath  and  finally  over  sulphuric  acid  in  a  desiccator. 

Place  about  1  g.  (Quant.)  of  cupric  sulphide  in  a  porcelain  boat, 
and  introduce  it  into  the  reduction  tube  used  in  Exp.  7b.  Pass  dry 

(72) 


Elective  Experiments.  73 

hydrogen  over  the  heated  sulphide  and  continue  the  experiment  until 
the  weight  of  the  residue  remains  constant.  Calculate  the  loss  in 
weight  of  the  sulphide,  and  the  quantities  of  cupric  and  cuprous  sul- 
phides that  are  equivalent  to  the  loss  of  32  g.  of  sulphur. 

Iron  pyrites  and  stannic  sulphide  also  yield  lower  sulphides  by  this 
method.  Illuminating  gas  may  be  substituted  for  hydrogen  as  a 
reducing  agent. 

50.  Sulphide    of    a    Metal    from    its    Sulphate.     (Quant.)     (Two 
students  working  together.)     a.     Weigh  about  1.5  g.  of  barium  sul- 
phate and  reduce  it  to  the  sulphide,  proceeding  as  in  the  previous 
experiment.     Heat  strongly  and  continue  the  reduction  till  constant 
weight  is  obtained.     Calculate  the  ratio  of  sulphide  to  sulphate. 

b.  Lead  and  zinc  sulphates  are  converted  into  sulphides  by  heating 
with  sulphur.  Weigh  about  1  to  2  g.  of  dry  lead  sulphate,  or  of  anhy- 
drous zinc  sulphate  in  a  crucible,  mix  with  an  equal  bulk  of  sulphur, 
and  then  heat  strongly  in  the  covered  crucible.  Repeat,  with  fresh 
additions  of  sulphur,  to  constant  weight.  Several  repititions  will 
probably  be  necessary.  Calculate  the  ratio  of  sulphide  to  sulphate. 

How  may  lead  sulphide  be  covered  into  sulphate?  Zinc  sulphide 
into  zinc  sulphate? 

51.  Volumetric  Composition  of  Ammonia. 

Fit  a  tube  at  least  60  cms.  long  and  closed  at  one  end,  with  a  one- 
hole  rubber  stopper  through  which  a  short  stemmed  drop  funnel  is 
passed,  in  such  a  way  that  the  end  of  the  stem  is  even  with  the  end  of 
the  cork.  Fill  the  dry  tube  with  chlorine.  Having  the  stop  cock  of 
the  drop  funnel  closed,  fill  the  short  stem  with  water  and  fix  the  cork 
in  the  tube.  Pour  into  the  funnel  about  5  cc.  of  ammonium  hydroxide, 
and  let  most  of  it  gradually  enter  the  tube.  When  the  reaction  is 
complete,  flow  in  dilute  sulphuric  acid  to  neutralize  the  excess  of 
ammonium  hydroxide  in  the  tube.  At  no  time  during  these  operations 
must  air  be  allowed  to  enter  through  the  funnel.  Finally  admit 
through  the  funnel  as  much  of  dilute  acid  or  water  as  will  enter  the 
tube.  Measure  the  volume  of  the  residual  gas  in  the  tube  and  compaer 
it  with  the  volume  of  the  tube. 

What  has  become  of  the  chlorine  that  filled  the  tube?     What  is 
the  residual  gas?     From  what  substance  did  it  come  and  what  vol- 
ume would  this  substance  occupy  in  the  free  state? 
Chem. — 6 


74  Laboratory  Experiments. 

52.  Qualitative  Separation  of  Lead,  Silver,  Mercurous,  and  Po- 
tassium Ions. 

a.  In  separate  tests   tube   place   about  10  cc.  of  the  following  so- 
lutions:      Lead    nitrate,    silver    nitrate,     mercurous    nitrate,     potas- 
sium nitrate.     Add  dilute  hydrochloric  acid  to  each  until  the  action 
is  complete.     (?)     Allow  the  precipitates  to  settle  and  then  pour  off 
the  supernatant  liquids  from  each  of  the  tubes.     Add  to  the  precipi- 
tates enough  cold  water  nearly  to  fill  the  tubes,  and  shake  the  con- 
tents.    After  the  precipitates  have  settled,  decant  the  liquids.     What 
compound  has  largely  been  removed  by  these  operations?     Try  next 
the  effect  of  hot  water  on  the  precipitates.     (?)     Try  the  effect  of 
ammonium  hydroxide  on  the  precipitates.     (?) 

b.  Mix  about  5  cc.  of  each  of  the  above  solutions  of  nitrates.     To 
the  mixture  add  dilute  hydrochloric  acid  until  precipitation  is  com- 
plete.    Of  what  does  the  precipitate  consist?     Filter  off  the  precipi- 
tate.    What  basic  ion  is  now  in  the  filtrate?     Wash  the  precipitate 
on  the  filter  with  a  little  cold  water,  and  reject  the  washings.     Next 
wash  the  precipitate  with  hot  water,  keeping  the  washings.     Divide 
the  filtrate  into  three  parts;  to  one  add  a  few  drops  of  dilute  sulphuric 
acid;  to  another  a  few  drops  of  potassium  chromate  solution,  and  to 
the  third  add  a  little  hydrogen  sulphide  water,  or  pass  in  a  few  bub- 
bles of  hydrogen  sulphide.     The  tests  are  taken  as  confirmatory  evi- 
dence for  the  presence  of  lead  ion  in  the  filtrate.     Write  the  equa- 
tions for  the  confirmatory  tests. 

Wash  the  precipitate  on  the  filter  with  hot  water  until  the  washings 
are  free  from  lead.  How  is  this  determined?  Reject  the  washings. 
Now  wash  the  precipitate  remaining  on  the  filter  with  ammonium 
hydroxide,  keeping  the  washings.  (?)  Acidify  the  filtrate  with  nitric 
acid.  (?)  What  separation  was  made  by  washing  the  precipitate 
with  ammonium  hydroxide?  Suggest  a  confirmatory  test  for  silver 
chloride. 

Dissolve  the  black  precipitate  on  the  filter  in  aqua  regia,  made  by 
adding  1  cc.  of  nitric  acid  to  3  cc.  of  hydrochloric  acid  and  warming. 
Dilute  with  water  the  solution  thus  obtained,  and  put  into  it  a  bright 
strip  of  copper.  After  several  minutes,  remove  the  strip,  and  wash 
and  rub  it.  (?) 

c.  Obtain  from  the  instructor  an  unknown  solution  that  may  con- 
tain one  or  more  of  the  above  ions.     Using  the  methods  in  b,  analyze 


Elective  Experiments.  75 

the  solution  for  lead,  silver,   mercury,  and  potassium,  recording  all 
steps  of  the  process,  even  those  giving  negative  results. 

53.  Analysis  of  a  Silver  Coin.  Quant.  (Two  students  working 
together.) 

Clean  a  dime,  weigh  it,  and  dissolve  it  in  diluted  nitric  acid,  by 
gentle  heating.  Evaporate  to  dryness  on  a  steam  bath,  and  then  re- 
dissolve  in  water.  Dilute  the  solution  and  make  it  up  to  250  cc. 
in  a  measuring  flask.  (Storeroom.)  Divide  the  solution  into  two 
portions,  one  of  100  cc.  and  the  other  of  150  cc.  Determine  the 
silver  in  the  smaller  sample  by  acidifying  it  with  a  few  drops  of 
nitric  acid,  warming  the  solution,  and  then  adding  dilute  hydrochloric 
acid  in  slight  excess,  a  drop  at  a  time  with  stirring.  Cover  the 
beaker  with  a  watch  glass  and  set  it  aside  in  a  dark  place  for  12 
hours.  Decant  the  liquid  through  a  weighed  Gooch  crucible  (Store- 
room) containing  an  asbestos  filter.  (Instructions.)  Transfer  the 
precipitate  to  the  filter,  wash  it  twice  with  hot  water,  dry  it, 
(Instructions.)  and  finally  weigh.  From  the  weight  of  silver  chloride 
calculate  the  silver  and  then  the  percentage  of  silver  in  the  coin. 
Save  the  filtrate. 

Precipitate  the  silver  in  the  other  sample  by  bringing  a  weighed 
piece  of  bright  sheet  copper  into  the  solution.  Stir  the  solution 
occasionally,  and  let  it  stand  till  a  drop  of  the  solution  brought  in 
contact  with  a  drop  of  hydrochloric  acid  shows  no  turbidity.  Then 
remove  the  copper,  rubbing  off  any  adhering  particles  of  silver  into 
the  beaker.  Decant  the  solution  through  a  small  quantitative 
filter  (Storeroom)  and  bring  the  silver  on  the  paper  by  inclining  the 
beaker  and  directing  a  fine  stream  of  water  on  it.  Wash  the  silver 
and  then  transfer  the  moist  paper  and  silver  to  a  weighed  crucible. 
Dry  over  a  low  flame  and  then  heat  till  the  paper  is  burned.  Use 
the  data  in  calculating  the  percentage  of  silver  in  the  coin.  From 
the  loss  in  weight  of  the  copper  calculate  its  equivalent  of  silver,  and 
and  quantity  of  copper  equivalent  to  107.9  g.  of  silver.  Save  the 
filtrate. 

Determine  the  copper  in  the  filtrates  by  depositing  it  electrolyt- 
ically  on  a  platinum  gauze  cathode.  For  details  of  the  procedure 
consult  Smith's  Electrochemical  Analysis. 

Read  up  about  silver  chloride  and  then  submit  to  the  instructor 
a  method  of  obtaining  free  silver  from  it. 


CHAPTER  XIV. 

INORGANIC  AND  ORGANIC  PREPARATIONS. 

54.     Silicon  Tetrafluoride. 

Arrange  an  apparatus  as  shown  in  Fig.  15.  Prepare  a  mixture 
of  15  g.  of  finely  ground  sand  and  10  g.  of  powdered  fluorspar.  Place 
the  mixture  in  a  flask  of  about  250  cc.  capacity,  and  pour  in  enough 
concentrated  sulphuric  acid  to  make  by  shaking  a  thin  paste  of  the 
mixture.  Connect  the  flask  with  a  dry  exit  tube  which  is  connected 


Fig.  15. 

with  a  long  delivery  tube  also  dry,  dipping  under  a  little  mercury 
contained  in  a  small  crucible,  standing  in  the  bottom  of  a  beaker  of 
about  200  cc.  capacity.  The  delivery  tube  and  mercury  must  be  dry> 

(76) 


Preparations.  77 

Nearly  fill  the  beaker  with  water  without  letting  any  of  it  enter  the 
delivery  tube.  Heat  the  flask  carefully,  and  note  the  formation  in 
the  beaker  of  the  jelly-like  silicic  acid.  Stir  occasionally,  and  con- 
tinue to  pass  the  gas  (what  is  it?)  until  the  entire  liquid  becomes  a 
jelly.  Disconnect  the  apparatus,  pour  the  contents  of  the  genera- 
tor into  the  waste-jar,  and  rinse  the  flask.  Remove  the  crucible  and 
mercury,  and  pour  the  mercury  into  the  bottle  for  impure  mercury. 
Do  not  pour  it  into  the  sink  or  waste-jar. 

Filter  the  contents  of  the  beaker  and  save  the  filtrate.  Wash  the 
jelly  twice  on  the  filter,  rejecting  the  washings.  Heat  a  portion  of  it 
with  a  small  volume  of  sodium  hydroxide  solution,  using  as  little  as 
is  necessary  to  dissolve  it.  Name  the  solution.  Add  hydrochloric 
acid  to  it  drop  by  drop.  What  is  the  precipitate  that  forms?  Dry 
the  rest  of  the  precipitate  by  pressing  it  between  filter  papers,  and 
then  heat  it  gently  until  a  white  powder  is  obtained.  (?) 

Test  the  filtrate  with  blue  litmus  paper.  (?)  To  one  portion  add 
a  solution  of  barium  chloride,  (?)  and  to  another  potassium  nitrate 
solution.  (?)  Shake  the  precipitates  and  examine  their  struc- 
ture. (?) 

Explain  the  chemical  changes  involved  in  every  step  of  the  experi- 
ment. 

55.  Silicon  Tetrachloride. 

Conduct  dry  chlorine  gas  over  pulverized  silicon  placed  in  a  hard 
glass  tube  which  is  heated  by  a  row  of  Bunsen  burners  or  in  a  com- 
bustion furnace.  Do  not  fill  the  tube  more  than  half  full  of  the  sil- 
icon. Collect  the  product  in  a  side  neck  flask  cooled  with  a  mixture 
of  ice  and  salt.  The  excess  of  chlorine  should  be  absorbed  by  caustic 
soda  solution  and  not  allowed  to  escape  in  the  room.  Rectify  the 
product  by  distilling  from  a  water  bath,  after  having  shaken  it  with 
a  little  mercury  to  remove  the  dissolved  chlorine  it  contains.  Note 
the  boiling  point  and  read  the  barometer.  Describe  the  product. 
Test  a  little  of  it  with  a  few  drops  of  water.  Write  the  reactions. 

56.  Purification  of  Sodium  Chloride. 

Prepare  about  150  cc.  of  a  cold,  saturated  solution  of  crude  salt 
by  grinding  the  salt  for  some  time  in  a  mortar  with  water.  Filter 
the  solution  into  a  beaker  and  pass  hydrogen  chloride  into  the  solu- 
tion. Prepare  the  gas  by  dropping  concentrated  sulphuric  acid  from 


78  Laboratory  Experiments. 


a  drop-funnel  on  common  salt  covered  with  concentrated  hydrochloric 
acid  in  the  ordinary  form  of  generator.  Deliver  the  gas  into  the 
solution  through  a  funnel  with  the  mouth  downward.  (?)  After 
considerable  precipitation  has  occurred  filter  off  the  salt  crystals. 
Wash  once  with  distilled  water,  and  remove  some  of  the  capillary 
water  by  means  of  a  suction  filter.  Transfer  the  salt  to  a  clean  evap- 
orating dish  and  dry  by  gentle  heating.  Reserve  about  5  g.  of  the 
product  for  some  tests  and  turn  over  the  balance  to  the  instructor. 
Compare  the  purity  of  the  product  with  the  original  sample  of 
salt  by  making  the  following  qualitative  tests: 

(a)  To  solutions  of  each,  add  sodium  carbonate  solution.     (?) 

(b)  Test  solutions  of  each  for  the  sulphate  ion. 

(c)  Test  solutions  of  each  for  the  magnesium  ion.     (K.  396;  S. 
430.)     (Instructions.) 

57.     Anhydrous  Ferric  Chloride. 

In  a  hard,  glass  tube  at  least  30  cm.  long,  place  several  coils  of 
bright,  iron  wire.  Fit  one  end  of  the  tube  with  a  one-hole  cork  and 
short,  glass  tube  and  mount  the  tube  for  heating  as  in  Exp.  7b.  Con- 
nect the  tube  with  an  apparatus  for  generating  dry  chlorine  (how  is 
the  gas  dried?)  and  push  the  iron  wire  toward  the  end  of  the  tube 
connected  with  the  generator.  Pass  a  gentle  stream  of  chlorine 
over  the  iron  and  heat  the  latter  moderately.  Vapors  of  ferric  chlo- 
ride will  soon  appear  and  condense  in  the  cool  portion  of  the  tube. 
Absorb  the  excess  chlorine  in  caustic  soda  solution.  After  some 
time  when  a  fair  quantity  of  sublimate  has  formed,  and  while  the  gas 
is  still  passing,  place  a  clean  dry  test  tube  over  the  open  end  of  the 
tube  and  sublime  the  ferric  chloride  into  the  test  tube  by  careful  ma- 
nipulation of  the  flame.  Preserve  the  crystalline  sample  of  ferric 
chloride  thus  obtained  by  sealing  off  the  tube  quickly  in  the  blast 
lamp. 

Test  small  portions  of  the  salt  remaining  in  the  tube  as  follows: 

(a)  Prepare  a  solution  of  the  salt  and  divide  it  into  four  portions. 
To  one  portion  add  ammonium  hydroxide  and  boil.     (?) 

(b)  Add  a  drop  of  the  solution  to  5  cc.  of  ammonium  thiocyanate 
solution.     (?) 

(c)  To  a  third  portion  add  a  few  drops  of  a  solution  of  potassium 
ferrocyanide.     (?) 


Preparations.  79 

(d)  To  the  fourth  portion  add  a  few  drops  of  concentrated  hydro- 
chloric acid  and  a  few  small  pieces  of  zinc.  From  time  to  time  test 
the  solution  by  pouring  a  drop  of  it  into  a  little  ammonium  thio- 
cyanate  solution.  (?) 

58.     Sulphuryl  Chloride.     (Two  students  working  together.) 

Chlorine  and  sulphur  dioxide  unite  when  exposed  to  sunlight  to 
form  a  liquid  known  as  sulphuryl  chloride.  When  chlorine  is  passed 
into  a  solution  of  camphor  in  sulphur  dioxide,  the  union  takes  place 
much  more  readily,  and  may  be  conducted  at  room  temperatures. 
The  camphor  remains  unchanged  in  the  process  and  so  exercises  a 
catalytic  effect. 

Set  up  a  generator  for  preparing  dry  chlorine,  and  another  one  for 
furnishing  sulphur  dioxide.  The  latter  may  however  be  conveniently 
taken  from  a  drum  of  liquid  sulphur  dioxide.  (See  instructor.)  Fit 
a  wide-mouth  bottle  with  a  three-hole  stopper  provided  with  two 
L-tubes,  whose  lower  ends  extend  nearly  to  the  bottom  of  the  bottle, 
in  which  the  synthesis  of  the  chloride  is  to  be  effected.  Place  about 
20  g.  of  camphor  in  the  bottle  and  set  it  in  a  dish  of  cold  water.  Now 
pass  in  slowly  dry  sulphur  dioxide  and  observe  that  the  camphor 
gradually  liquifies.  What  is  the  boiling  point  of  liquid  sulphur  diox- 
ide? How  do  you  explain  its  condensation  at  the  higher  temperature? 
When  most  of  the  camphor  has  dissolved  start  the  stream  of  chlorine 
into  the  liquid,  and  continue  passing  both  gases  at  about  the  same  rate 
until  the  volume  of  the  solution  has  increased  about  25  cc.  At  least 
an  hour  will  be  required  for  this  operation.  Toward  the  close  of 
the  experiment  discontinue  the  flow  of  sulphur  dioxide,  but  main- 
tain the  stream  of  chlorine.  (Object  of  this?)  Before  setting  up  the 
chlorine  generator  calculate  the  quantities  of  the  materials  to  be  used 
for  preparing  the  chlorine  for  40  g.  of  sulphuryl  chloride. 

Transfer  the  liquid  to  a  dry  distilling  flask,  connect  to  a  condenser, 
and  distill  on  a  water-bath,  collecting  the  distillate  in  a  dry  receiver. 
Finally  rectify  the  distillate  by  a  second  distillation  (Exp.  27),  noting 
the  temperature  at  which  the  liquid  distils,  and  read  the  barometer. 
Weigh  the  final  product,  and  estimate  the  cost  of  the  chemicals  used 
for  its  preparation,  assuming  that  all  of  the  by-products  are  waste 
materials. 

Pour  a  few  drops  of  water  into  a  small  portion  of  the  liquid.  (?) 
Test  the  liquid  resulting  from  the  interaction  for  the  sulphate  ion. 


80  Laboratory  Experiments. 

59.  Precipitated  Chalk. 

With  the  co-operation  of  the  instructor  secure  about  500  cc.  of  the 
spent  liquors  from  one  of  the  Kipp  generators  used  for  the  supply  of 
carbon  dioxide.  Make  the  solution  slightly  alkaline  by  adding  am- 
monium hydroxide,  and  then  heat  it  to  boiling.  What  substances 
are  probably  present  in  the  precipitate?  Filter  off  the  precipitate 
and  reject  it.  The  nitrate  should  be  clear  and  colorless.  Take  about 
100  cc.  of  the  nitrate,  cool  it  to  20°,  and  determine  its  density  with 
the  hydrometer,  and  by  referring  to  the  table  on  p.  331  of  the  phys- 
ical-chemical tables  of  Landolt,  Bo'rnstein,  and  Meyerhoffer  deter- 
mine the  approximate  quantity  of  calcium  chloride  in  the  entire  solu- 
tion you  are  using  in  the  experiment.  Then  calculate  the  quantity 
of  ammonium  carbonate  required  to  precipitate  the  calcium  as  car- 
bonate. Dissolve  this  amount  of  ammonium  carbonate  in  sufficient 
cold  water  to  make  a  10  per  cent  solution.  Dilute  the  nitrate  with 
distilled  water  till  the  volume  is  about  2-3  liters,  and  heat  it  to  boil- 
ing in  a  large  avaporating  dish.  Now  gradually  stir  in  the  ammo- 
nium carbonate  solution.  How  will  you  ascertain  whether  you  have 
added  enough?  Allow  the  precipitate  to  settle,  and  then  decant  off 
the  supernatant  liquid  and  set  it  aside.  What  does  it  contain?  Wash 
the  precipitate  with  water  made  slightly  alkaline  with  ammonium 
hydroxide  until  the  chlorides  are  washed  out.  (?)  Finally  drain 
the  precipitate  well,  dry  it,  and  weigh  it.  How  does  the  yield  com- 
pare with  the  quantity  of  carbonate  as  calculated  from  the  concen- 
tration of  calcium  chloride  in  the  original  solution?  Consult  a  price- 
list  of  the  chemical  reagents  and  make  an  estimate  of  the  compara- 
tive cost  of  the  product  that  you  have  prepared. 

Recover  the  salt  the  nitrate  contains. 

What  is  the  chief  use  of  precipitated  chalk? 

60.  Chrome  Alum. 

Dissolve  10  g.  of  potassium  dichromate  in  water,  and  add  the 
amount  of  sulphuric  acid  necessary  to  form  potassium  sulphate  and 
chromium  sulphate.  Warm  the  solution  and  then  add  gradually 
enough  alcohol  to  make  the  solution  bright  green.  Ten  cc.  of  95 
per  cent  alcohol  will  answer.  Note  the  odor  of  the  vapors  that  are 
produced.  (?)  (S.  478.)  Set  the  greater  portion  of  the  solution 
aside  to  evaporate  spontaneously  and  hang  a  string  in  the  solution. 
To  obtain  a  large  crystal  with  all  faces  well  formed  its  growth  should 


Preparations.  81 


not  be  checked  in  certain  directions  by  groups  of  other  crystals  clus- 
tering about  it.  Nurse  several  of  the  crystals  that  have  formed  on 
the  string  by  removing  the  intervening  ones:  If  any  crystals  have 
formed  on  the  bottom  or  sides  of  the  vessel,  transfer  the  solution  to 
another  vessel  and  filter  it.  (?)  Again  hang  the  string  in  the  solu- 
tion, and  in  this  way  secure  several  good  sized  and  symmetrically 
developed  crystals. 

Concentrate  the  smaller  portion  of  the  solution  on  the  water-bath 
till  crystals  appear.  Examine  the  form  and  color  of  the  crystals 
from  both  portions.  (?)  What  is  the  color  of  their  solution  in 
water?  (?) 

Suggest  a  method  of  converting  the  chrome  alum  back  into  potas- 
sium dichromate. 

61.  Oxalic  Acid. 

Place  50  grams  of  cane  sugar  in  a  flask  of  about  1  liter  capacity  and 
gradually  add  300  cc.  of  nitric  acid  of  sp.  gr.  1.38.  Warm  gently, 
if  necessary,  to  start  the  reaction,  but  as  soon  as  the  evolution  of  the 
red  fumes  begins  remove  the  source  of  heat.  Allow  the  action  to 
proceed  under  the  hood  till  no  more  red  fumes  are  formed.  In  an 
evaporating  dish  boil  down  the  liquid  to  about  one-sixth  of  its  bulk. 
Cool  the  liquid  and  collect  the  crystals  on  a  funnel,  draining  them 
well.  Press  them  between  filter  paper  and  recrystallize  them  from 
a  little  hot  water.  Concentrate  the  mother  liquor  further  to  secure 
a  second  drop  of  crystals.  Weigh  the  product.  Write  the  reactions. 
Compute  the  percentage  yield.  Describe  the  substance  you  have 
prepared.  Dissolve  a  little  of  it  in  water,  add  a  drop  of  ink.  Ex- 
plain. Heat  a  crystal  of  oxalic  acid  on  platinum  foil.  Do  the  crys- 
tals contain  water  of  crystallization? 

62.  Preparation  of  lodoform  from  Acetone. 

To  100  cc.  of  about  a  20  per  cent  sodium  carbonate  solution  in  a 
beaker  of  300  to  400  cc  capacity  add  20  cc.  of  acetone  and  warm  the 
whole  to  70°  C.,  preferably  on  a  water  bath.  Pulverize  10  grams 
of  iodine  thoroughly  in  a  mortar  and  add  this  to  the  liquid  in  very 
small  amounts  at  a  time,  stirring  constantly.  When  the  dark  brown 
color  has  disappeared,  allow  to  cool.  Filter  off  the  crystals  of  iodoform ; 
wash  them  thoroughly  with  cold  water,  and  then  dry  them  with  filter 
paper  by  pressing  the  crystals  between  folds  of  the  latter.  Weigh  the 


82  Laboratory  Experiments. 

•product.  Write  the  reaction.  How  much  iodoform  ought  you  to 
have  obtained  from  the  10  grams  of  iodine  used.  The  amount  you 
have  obtained  divided  by  the  amount  that  theoretically  might  be 
obtained  is  the  percentage  yield.  Compute  the  percentage  yield  in 
this  case.  What  is  iodoform  used  for?  Why  is  it  costly? 

63.     Soap    and     Glycerine.     (Two    students    working    together.) 

Look  up  the  composition  of  lard  or  tallow  and  the  reaction  involved 
in  the  preparation  of  soap  by  the  interaction  of  sodium  hydroxide  and 
grease.  Write  the  reaction  and  calculate  the  theoretical  quantity 
of  grease  that  will  react  with  25  g.  of  sodium  hydroxide.  In  practice 
a  large  excess  of  lye  is  required  to  make  the  reaction  successful;  there- 
for use  less  than  half  the  theoretical  quantity  of  grease.  Dilute  the 
solution  of  sodium  hydroxide  to  about  300  cc;  Pour  half  of  it  into  an 
iron  pan  having  a  capacity  of  at  least  a  liter;  add  about  150  cc.  of  water 
-and  the  grease.  Boil  this  for  half  an  hour,  adding  hot  water  from 
time  to  time  so  as  to  maintain  the  original  volume.  Now  add  the  rest 
•of  the  lye  and  boil  for  an  hour  or  longer,  supplying  the  water  lost  by 
evaporation.  Concentrate  it  by  boiling  off  about  a  third  of  the  water. 
Disslove  50  g.  of  common  salt  in  the  solution  before  the  boiling  is 
stopped. 

Allow  to  cool  until  the  next  laboratory  period.  The  soap  will  rise 
to  the  top  and  solidify,  while  the  glycerin,  excess  of  lye  and  salt  will 
be  in  the  solution. 

Make  the  solution  left  from  the  soap  slightly  acid  with  hydrochloric 
acid;  filter  if  necessary.  Evaporate  the  filtrate  in  a  porcelain  dish 
on  a  water  bath  to  as  near  dryness  as  possible.  Extract  the  glycerine 
from  the  residue  with  alcohol  and  evaporate  off  the  alcohol  on  the 
water  bath.  Describe  the  taste,  color,  consistency,  etc.,  of  the  product, 
and  compare  it  with  the  laboratory  specimen.  What  use  is  made  of 
the  large  quantities  of  crude  glycerine  obtained  as  a  by-product  in  the 
manufacture  of  soap?  Why  is  glycerine  regarded  as  a  kind  of  alcohol? 
Compare  the  composition  of  the  glycerin  molecule  with  that  of  calcium 
hydroxide  and  sodium  hydroxide.  What  is  the  essential  difference? 
Compare  the  composition  of  fat  with  that  of  soap.  What  is  the 
essential  difference?  Is  fat  a  salt? 

Prepare  a  little  of  the  acid  from  which  soap  may  be  regarded  as 
derived  by  dissolving  completely  about  10  g.  of  fine  shavings  of  a  good 


Preparations.  83 

grade  of  laundry  soap  in  hot  water.  Make  the  hot  solution  acid  with 
dilute  sulphuric^'acid.  (?)  Boil  slowly  and  watch  for  an  oily  layer 
of  fatty  acid.  Set  the  liquid  aside  to  cool.  If  the  acid  layer  solidifies 
remove  it  and  treat  it  with  boiling  water  to  wash  out  any  sulphuric  acid 
which  may  have  become  incorporated  with  it.  Solidify  as  before 
and  examine  its  texture,  color,  odor,  etc. 


APPENDIX. 

I.     ABBREVIATIONS. 

The  following  abbreviations  are  used  in  this  outline: — 
cc.  cubic  centimeter  (s)  mg.  milligram  (s) 

cm.  centimeter  (s)  mm.     millimeter  (s) 

g.  gram  (s)  1.  liter  (s) 

kg.  kilogram  (s)  Quant,  quantitative 

K.  Kahlenberger's  Outlines  of  Chemistry. 

S.  Alexander  Smith's  General  Chemistry  for  Colleges. 

II.     Measures  of  Length  and  Equivalents. 

1  meter  =  10  decimeters  =  100  cm.  =  1000  mm.  =  39.37  inches. 
1  foot  =  30.48  cm.;  1  in.  =  2.54  cm.  =  25.4  mm. 

m.     Measures    of    Capacity. 
1  liter  =  1000  cc.  =  1.057  quarts. 
1  quart  (U.  S.,  liquid)  =  946.36  cc. 
1  ounce,  avoirdupois  =  28.35  cc. 
1  pound,  avoirdupois  =  453.6  cc. 

IV.     Measures  of  Weight. 

1  kg.  =  1000  g.  =  weight  of  1000  cc.  of  water  at  4°  =  2.205  Ibs. 
1  oz  (av.)       =  28.35  g.;  1  g.  =  15.43  grains;  1  mg.  =  0.001  g. 

V.    Correction    of    Barometric    Readings. 

To  reduce  the  reading  taken  at  room  temperatures  to  the  correspond- 
ing height  of  a  column  of  mercury  at  0°,  subtract  the  proper  number  in 
the  correction  column  from  the  actual  reading  in  millimeters. 


Temp. 
12 

Corr'n. 
1.6 

Temp. 
17 

Corr'n. 
2.2 

Temp. 
22 

Corr'n. 
2.85 

Temp. 
27 

Corr'n- 
3.5 

13 

1.7 

18 

2.3 

23 

3.0 

28 

3.6 

14 

1.8 

19 

2.5 

24 

3.1 

29 

3.75 

15 

2.0 

20 

2.6 

25 

3.2 

33 

4.15 

16 

2.1 

21 

2.7 

26 

3.35 

35 

4.55 

(84) 


Appendix.  85 

VI.    Pressure  of  Water  Vapor,  or  Aqueous  Tension,  in  mm.  of  Mer- 
cury. 


Temp. 
0° 

Press. 
4.6 

Temp. 
19 

Press. 
16.4 

Temp. 
29 

Press. 
29.8 

5 

6.5 

20 

17.4 

30 

31.55 

10 

9.2 

21 

18.5 

31 

33.4 

12 

10.5 

22 

19.7 

32 

35.4 

13 

11.2 

23 

20.9 

33 

37.4 

14 

11.9 

24 

22.2 

34 

39.6 

15 

12.7 

25 

23.55 

35 

41.85 

16 

13.55 

26 

25.0 

40 

55.0 

17 

14.45 

27 

26.5 

60 

149.2 

18 

15.4 

28 

28.1 

100 

760.0 

VII.     Reduction    of    a    Gas    Volume    to    Standard    Conditions. 

Assuming  a  volume  of  air  standing  over  water  measures  400  cc.  when 
the  barometer  reads  742  mm.  and  the  temperature  is  27°  C.,  what  would 
be  the  volume  of  the  air  at  standatd  conditions,  0°  C.  and  760  mm.? 

If  the  temperature  near  the  barometer  is  27°,  the  column  of  mercury 
is  longer  than  it  would  be  at  0°.  The  correction  to  be  deducted  ap- 
pears in  Appendix  5  as  3.5  mm.  Hence  the  corrected  barometric 
reading  is  742-3.5  or  738.5  mm. 

The  gas  volume  measured  is  not  dry  air,  but  a  mixture  of  air  and 
water  vapor,  whose  combined  pressure  is  738.5  mm.  By  referring  to 
Appendix  6,  we  see  that  the  partial  pressure  of  the  water  vapor  at 
27°  is  26.5  mm.  Hence  the  partial  pressure  of  the  air  is  738.5-26.5 
or  712  mm.,  which  is  equivalent  to  saying  that  the  air  when  dry  and 
in  a  space  of  400  cc.  would  exert  a  pressure  of  712  mm. 

What  would  be  the  volume  of  the  air  if  the  pressure  were  increased 
to  760  mm.?  The  increase  in  pressure  compresses  the  gas,  and  while 
the  temperature  is  kept  constant,  the  volume  varies  inversely  as  the 
pressure.  (Boyle's  law.)  The  volume  will  then  change  in  the  ratio 
of  712  to  760,  or  the  volume  under  a  pressure  of  760  mm.  and  27°  is 
given  by  400  x  -ffg- 


86  Appendix. 

Now  lowering  the  temperature  to  0°,  another  contraction  in  volume 
occurs.  Since  the  volume  of  a  gas  varies  directly  as  the  absolute 
temperature,  the  pressure  remaining  constant,  (Law  of  Charles)  and 
the  temperatures  here  considered  are  273°  absolute  (  =  0°  C.)  and  300° 
Abs.  (  =  27°  C.),  the  reduction  factor  is  f££  The  complete 
equation  for  reduction  of  gas  to  standard  conditions  then  becomes 

400  X      fU     X      iU     =     354.9  cc. 

The  factors  in  this  equation  are  the  observed  volume,  V,  the  partial 
pressure,  P,  the  standard  pressume  P0,  the  standard  temperature, 
To  the  observed  temperature,  T,  and  the  volume  reduced  to  standard 
conditions,  Vo.  We  may  then  write  the  general  equation  for  the  re- 
duction of  a  gas  to  standard  conditions  in  the  form, — 

V  P  T 
V0=  ^~or  VPT0=V0P0T 


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